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INTRODUCTION TO GENERAL 
CHEMISTRY 



INTRODUCTION 

TO 

GENERAL CHEMISTRY 



HERBERT N. McCOY 

and 

ETHEL M. TERRY 



CHICAGO, ILLINOIS 
1919 



f 






Copyright igig By 
Herbert N. McCoy and Ethel M. Terry 



All Rights Reserved 



Printed August 191 9 



Composed and Printed By 

The University of Chicago Press 

Chicago, Illinois, U.S.A. 



AUG ' 

©CU530526 






PREFACE 

This text has been in process of growth at the University 
of Chicago since 19 13. At that time a synopsis of the first 
nine chapters was printed. This was followed, in 1916-1 7, by the 
completed work of the first fifteen chapters. Chapters xvi-xix 
were added and used in class work during 19 18 and 19 19. 

The book has been written for college Freshmen, and, as its 
title implies, it is intended to serve as an introduction to general 
chemistry. In consequence we have aimed to present a contin- 
uous and connected story in teachable form and have not 
attempted to give extensive descriptive and numerical data 
where such matter is of little interest to the student or is not 
needed for the development of important principles. 

Inasmuch as the choice and arrangement of topics in the 
earlier part of this book depart noticeably from the familiar 
order, some explanation seems necessary. We shall therefore 
sketch briefly the plan of the less conventional chapters, 
together with the " philosophy of arrangement" which has 
resulted in the scheme presented. 

The first chapter, which is brief, deals with the measurement 
of gases and the gas laws. In the next five chapters the most 
fundamental concepts of the science of chemistry are developed. 
These include: indestructibility of matter, idea of a pure sub- 
stance, decomposition of pure substances, elements, analysis of 
substances and percentage composition, the law of definite 
composition, derivation of formulae. 

Chapter v shows how chemical formulae are derived from a 
knowledge of percentage composition and gas or vapor density. 
This development keeps as close as possible to the arguments 
of Avogadro and Cannizzaro and shows how formulae are ob- 
tained by methods independent of the atomic-molecular hypoth- 
esis. Reference, at this stage, to combining weights and 
chemical equivalents is purposely avoided, for the reason that 
the history of chemistry between the time of Avogadro and 



vi Preface 

that of the epoch-making paper of Cannizzaro (1858) shows 
the fallacy of trying to develop formulae and fix atomic 
(symbol) weights by any method other than that proposed by 
Avogadro and elucidated by Cannizzaro. The sixth chapter 
introduces the use of equations before the atomic-molecular 
hypothesis is studied. This plan has the great advantage of 
fixing in the student's mind the fundamental relationship 
between equations and the quantitative experimental data 
such equations represent. 

Chapters ii to vi inclusive form a compact division of the 
subject, in which the argument, illustrated at every step by 
experimental data, is substantially continuous. In these chap- 
ters we have aimed at a logical development of the subject 
without the introduction of any matter that does not serve to 
illustrate the topics under discussion. 

The next three chapters, vii, " Acids, Bases, and Salts — I"; 
viii, "Water and Solutions"; and ix, "Acids, Bases, and Salts — 
II," are introduced at this point for very definite reasons. In 
the first place, it is obvious to every teacher that much of the 
beginner's w r ork will deal with acids, bases, and salts and their 
solutions. It is our opinion that a knowledge of these topics 
is best obtained by studying them directly and specifically, 
both in the classroom and in the laboratory. Our plan provides 
for laboratory work by the student, following closely the content 
of these three chapters. This laboratory work is interesting to 
students, since they like to make and crystallize a variety of 
salts. It also gives good training in technique and is not 
difficult either experimentally or theoretically, while at the 
same time it offers a wealth of material for practice in writing 
equations and solving problems. The most important reason, 
however, for the introduction of the early study of acids, 
bases, and salts is to supply the indispensable data needed later 
for the understanding of the ionic hypothesis. 

Chapters x and xi present the kinetic-molecular and atomic 
hypotheses respectively. It will be noted that these subjects 
follow the development and use of formulae instead of preceding 
them. This emphasizes the generally overlooked fact that for- 



Preface vii 

mulae are in no way, of necessity, dependent upon the molecular- 
atomic hypothesis. At the same time the student is in a position 
to appreciate more readily the molecular-atomic hypothesis be- 
cause it furnishes a plausible explanation of facts already familiar 
to him. 

Since the early chapters of the text contain as much infor- 
mation about hydrogen and oxygen as the student needs, or 
can fully appreciate at the start, the formal discussion of these 
elements has been postponed to chapter xiv. Chapters xiii 
("Chemical Equilibrium"), xv (" Oxidation and Reduction"), 
and xvi ("Heat and Energy") present important theoretical 
matters. 

The extended discussion given of the ionic hypothesis needs 
no apology. In this connection we have introduced a new 
method of graphic representation of ionic equilibrium. In 
chapter xx, on "Electrochemistry," the electronic conception 
of reactions, including oxidation and reduction, is discussed. 

The arrangement from this point on needs little comment. 
The authors recognize, of course, that organic chemistry (xxv 
and xxvi), "Theory of Dilute Solutions" (xxvii), "Disperse 
Systems" (xxviii), and "Radioactivity" (xxxii) are optional 
studies for beginning general chemistry courses. We have 
included these subjects, since the interest which they have 
aroused has apparently justified so doing. 

The 880 sections of the book have been not only titled but 
numbered, in order that frequent cross-references might be 
given. The student is thus constantly informed of the mate- 
rial on which each discussion is based. Instructors will find the 
references helpful in laying out work if they desire to skip 
certain portions of the text, for instance with classes of students 
who have had a secondary-school course in chemistry. 

The great importance of close connection between the work 
of the laboratory and that in class has been kept constantly 
in mind in planning this text. A laboratory guide (.4 Labora- 
tory Outline for General Chemistry, by McCoy and Terry), which 
follows strictly the arrangement of the text, has been written 
to accompany the latter. 



viii Preface 

In addition a pamphlet has been prepared containing those 
details of the lecture experiments which are important to the 
lecture assistant but of no interest to the elementary student. 
This pamphlet will be available for teachers using the text. 

We are indebted to Professor W. D. Harkins, of the Uni- 
versity of Chicago, for the contribution of sections 6 and 7 
(pp. 3 and 4) of the text and for the use of Fig. 117; to Mr: 
Leo Finkelstein for the drawings of Figs. 1 to 43; and to Dr. 
R. D. Mullenix for the seventy-seven additional drawings, as 
well as for valuable criticism. 

The Authors 

Chicago, III. 
July, 1919 



ERRATA 

Page 39, line 20, change members to numbers. 

Page 62, line 28, change to to of. 

Page 71, line 31, change 100 to 540. 

Page 79, Fig. 27, change PbN0 3 to Pb(N0 3 ) 2 . 

Page 79, Fig. 27, change BaCl 2 to BaCl 2 '2H 2 0. 

Page 128, line 2, after element add in the formula of the substance. 

Page 149, line 15, change bromine to bromide. 

Page 150, line 18, change AlBr to AlBr 3 . 

Page 151, line 2, change bromine to bromide. 

Page 211, lines 14 and 15, change hypochlorate to hypochlorite. 

Page 213, line 2, change HC0 4 to HC10 4 . 

Page 296, line 17, change Storey to Stoney. 



TABLE OF CONTENTS 

CHAPTER PAGE 

I. Introduction — Laws of Gases i 

II. The Burning or Substances — Oxygen 6 

III. Pure Substances — Elements 16 

IV. The Law of Definite Composition 26 

V. Symbols and Chemical Formulae 40 

VI. Chemical Equations 50 

VII. Acids, Bases, and Salts — I 57 

VIII. Water and Solutions 68 

IX. Acids, Bases, and Salts — II ........ 86 

X. The Kinetic Theory of Matter and the Molecular 

Hypothesis 109 

XL The Atomic Hypothesis and Atomic Weights . . . 121 

XII. The Halogens and Their Compounds with Hydrogen 

and Metals 135 

XIII. Chemical Equilibrium 158 

XIV. Hydrogen and Oxygen 176 

XV. Oxidation and Reduction . 194 

XVI. Heat and Energy 214 

XVII. The Ionic Hypothesis 227 

XVIII. Applications of the Ionic Hypothesis 251 

XIX. Applications of the Ionic Hypothesis. Reactions 

Involving Changes of State 275 

XX. Electrochemistry 296 

XXL Nitrogen and Ammonia 323 

XXII. Nitric Acid and the Oxides of Nitrogen .... 341 

XXIII. Phosphorus and Its Compounds 363 

XXIV. Sulfur and Its Compounds 375 

XXV. Carbon and Carbon Compounds, Organic Compounds— 

II 393 

ix 



x Table of Contents 

CHAPTER PAGE 

XXVI. Organic Compounds — II 426 

XXVII. Theory of Dilute Solutions 452 

XXVIII. Disperse Systems 470 

XXIX. The Atmosphere and Related Topics 491 

XXX. Some Additional Elements and Their Compounds . 514 

XXXI. Classification or the Elements. The Periodic System 539 

XXXII. Radioactivity and the Nature of Matter. . . . 568 

Index 593 



CHAPTER I 
INTRODUCTION— LAWS OF GASES 

i. A Knowledge of Physics Prerequisite for Chemistry. — 

The sciences of physics and chemistry are so closely related that 
the latter may be considered an extension of the former. A 
knowledge of physics is therefore necessary for an adequate 
understanding of chemistry, and it is to be assumed that the 
student taking up chemistry has had at least a one-year high- 
school course in physics. 

2. The Three Forms of Matter: Gases. — In his work in 
physics, the student will have learned the meaning of the term 
matter, which may be defined as anything which occupies space 
and has weight. He will have learned, also, that matter may 
exist in three forms: solid, liquid, and gaseous. Since gases are 
less tangible than solids and liquids, we shall first take up the 
study of air, the most familiar of all gases. That 

air has the two attributes just mentioned as belong- 
ing to all forms of matter may readily be shown by 
experiment. 

3. Air Occupies Space and Has Weight. — If a 
drinking glass or .beaker be thrust, mouth down- 
ward, into a vessel of water, the water does not 
enter until the glass is tilted to allow the air to 
escape. This shows that air occupies space. 

That air has weight may be shown by weighing 
a flask, first empty and afterward filled with air. 
The flask (Fig. 1) should be round-bottomed and yig. i 

have a capacity of 250 to 500 c.c. It is fitted with 
a tight rubber stopper carrying a glass stopcock. The air is 
first pumped out by means of an efficient air pump; the stop- 
cock is then closed and the flask counterbalanced with weights. 
When the stopcock is opened the inrush of air can be heard, 
and it is easy to observe that there is an appreciable increase 




2 Introduction to General Chemistry 

in weight. Since air occupies space and has weight, it is un- 
doubtedly a form of matter. 

One liter of air weighs more than a gram and the air contained 
in a room 12 feet square and 12 feet high would weigh about 100 
pounds. At the earth's surface air exerts a pressure of about 
15 pounds on every square inch of surface. The existence of 
this pressure may readily be shown by means of the following 
experiment. A tin can with a narrow neck (such as is often used 
for shipping alcohol, etc.) and of about 1 gallon capacity is 
fitted with a stopper carrying a glass tube, by means of which 
the air rilling the can may be pumped out. Usually, before the 
exhaustion of the air is complete, the can is crushed by the 
pressure of the air on the outside — a pressure which is now 
no longer balanced by the equal and opposite pressure on the 
inside. 

4. The Effect of Pressure on Volume: Boyle's Law. — The 
atmospheric pressure is measured by means of the barometer. 
At the sea-level the normal barometric pressure serves to support 
a column of mercury 76 cm. high. The effect of pressure upon 
the volume of air was first studied by Robert Boyle in the seven- 
teenth century. Boyle found that the volume of a given portion 
of air was inversely proportional to the pressure. This relation 
is known as Boyle's law. If we represent the pressure by P and 
the volume by V, then PV = & constant. 

5. The Effect of Temperature on Volume: The Law of 
Charles. — In scientific work we use the Centigrade thermometer, 
the scale of which is so constructed that the freezing-point of 
water is o°, while the boiling-point is ioo°. The effect of tempera- 
ture upon the volume of a given portion of air at a fixed pressure 
was studied over a century ago by Charles and by Gay Lussac. 
It was found that the volume of the air increased 1/273 of its 
volume at zero for each increase of i° C. This statement is 
known as the law of Charles, or sometimes also as the law of 
Gay Lussac. 

6. The Gas Thermometer: Absolute Temperature. — An 
experiment will show that if 273 c.c. of air contained in a flask 
or cylinder at o° C. is heated to ioo° C. the volume will change to 



Introduction — Laws of Gases 3 

373 c.c. Such an apparatus is called an air thermometer and 
temperatures may be measured in this way instead of by 
the expansion of mercury, as in ordinary thermometers. At 
i° C. the volume of the air is 274 c.c; at 2 it equals 275 c.c. 
and thus the volumes in the following table correspond to 
the temperatures given. 



TABLE I 



Volume 


Degrees 


Volume 


Degrees 


in c.c. 


Centigrade 


in c.c. 


Centigrade 


373 minus 


273 = 100 


293 minus 


> 273= 20 


372 " 


" = 99 


283 " 


" = IO 


363 " 


" = 90 


273 " 


" = O 


323 


" = 50 


263 


" =-10 


3i3 " 


" = 40 


253 " 


" =—20 


3°3 " 


" = 30 


243 


" =-30 



Since the zero of the Centigrade thermometer is arbitrarily 
chosen, being the temperature of the freezing of water (the stu- 
dent is already familiar with the Fahrenheit zero, which is at a 
lower temperature), it would be possible and convenient to use 
a temperature scale in which the volumes of the air in the air 
thermometer as described are taken as the temperatures. Since 
temperatures on the Centigrade scale are obtained by subtracting 
273 from the corresponding air thermometer temperatures, the 
zero of the air thermometer or gas scale must be 273 degrees lower 
than the Centigrade zero, or 273 degrees below the freezing-point. 
These air-thermometer temperatures are usually called the abso- 
lute temperatures; the absolute temperature may therefore be 
defined as the Centigrade temperature plus 273 degrees. Since 
most other gases act like air they may be used in the gas ther- 
mometer, and it is evident that if a certain amount of gas is used 
in such an experiment, no matter what its volume may be at the 
freezing-point of water, the volume will always vary with the 
temperature in the same ratio as the absolute temperature, pro- 
vided the pressure on the gas is kept constant. 

7. Problems. — We may now consider a few simple problems 
based on the two laws of gases just discussed. 



4 Introduction to General Chemistry 

Problem i: The volume of a certain amount of air at 2 7 C. 

is 1,000 c.c. What would its volume be at 12 7 C. if the pressure 

is kept constant? 

Centigrade temperature+ 2 73 = absolute temperature 

27°+ 2 73 = 3 00 
1 27°+ 273 = 400 

The volume of the gas must therefore increase in the ratio of 
400 to 300, or it will become 

ioooc.c.X— - = i333 -3 c - c - 
300 

Problem 2: Let the original pressure on the gas in Problem 1 
be 60 cm. of mercury (or -7 of the ordinary pressure of the atmos- 
phere). What will be the final volume of the gas if the pressure 
is increased to 100 cm. of mercury? An increase of pressure 
must decrease the volume of the gas, and in the ratio of the 

,_ 60 
pressures, 60 to 100, or by . 

a) Let the change of pressure come after the change of 
temperature as given in Problem 1 : 

then 

1333 .3 c.cX = 800 c.c. (final volume) Ans. 

100 

b) Let the change of pressure take place first : 

1000 c.cX — = 600 c.c, volume after the pressure change. 

The temperature change would then change the volume as 
follows: 

60c c.c. X — = 800 c.c. (final volume) Ans. 
300 

It is thus seen that the same answer is obtained, no matter 
which step in the problem is worked first, so the whole problem, 
1 and 2 together, may be stated in one expression as follows : 

400 60 
ioooc.c.X — X — = 800 c.c. Ans. 
300 100 



Introduction — Laws of Gases 5 

Problem j: Suppose that 1,000 c.c. of air at 20 C. and 70 cm. 
pressure is cooled to o° and that at the same time the pressure 
is increased to 76 cm. Find the final volume. 

When a gas is at the temperature of o° C. and under a pressure 
of 76 cm. (the normal atmospheric pressure at sea-level) it is said 
to be at standard conditions. 

Problem 4: Find the volume at standard conditions of 400 c.c. 
of air measured at 30 and 72 cm. 

8. Steam Is the Gaseous Form of Water. — It is well known 
that when water is heated it passes into steam. The white cloud 
which is frequently spoken of as steam is not really steam, but 
is composed of minute droplets of water. If we boil water in a 
glass flask the space above the water is filled with steam, but 
we notice that the steam is entirely invisible and that the visible 
cloud forms only when the steam cools and condenses to liquid 
droplets. Water in the form of steam is, like air, a gas. When 
we boil any liquid like alcohol or mercury the liquid passes into 
the state of a gas or vapor, as it is sometimes also called. The 
gas or vapor when cooled condenses to the liquid form of the sub- 
stance. 

9. Change of Form of Matter with Change of Temperature. — 
Just as water when cooled solidifies to ice, so every other liquid 
substance solidifies when sufficiently cooled. We speak of steam 
and ice as the gaseous and solid forms respectively of water. 
The substance known as moth-balls is called naphthalene by the 
chemist; it is a solid at ordinary temperatures, but when heated 
it melts to a colorless liquid, and when heated still hotter it 
boils, giving a colorless vapor, which is naphthalene in the form 
of a gas. When this gas is cooled it condenses to a liquid, which 
when cooled still further solidifies or freezes, giving solid naph- 
thalene again. Behavior like that of water and naphthalene 
is met with in the case of very many other substances. They 
can exist in three different forms, gas, liquid, and solid, according 
to the temperature. 



CHAPTER II 



THE BURNING OF SUBSTANCES— OXYGEN 

io. Burning Substances Require Air. — The history of chem- 
istry shows that the discovery of the real nature of the process of 
burning was one of the most important, if not the most important, 
in the development of the whole science. That air is needed for 
the burning of a substance is, in general, well known, and can 

easily be shown by many simple 
experiments. For example, if we 
place an inverted drinking-glass 
Q I over a burning candle standing on 

a table (Fig. 2), the flame quickly 
grows smaller and smaller and soon 
goes out, the glass having cut off 
l the needed supply of air. 
Fig. 2 A still more interesting and in- 

structive experiment may be made 
with phosphorus, a substance which burns very readily in the 
air, giving off clouds of white 
smoke. A piece of phosphorus of 
the size of a pea is placed on a 
cork floating on water and covered 
with a bell-jar (Fig. 3). When a 
heated wire passing through the 
tight-fitting stopper of the jar is 
brought in contact with the phos- 
phorus the latter takes fire and 
burns with the production of light 
and heat and the formation of a 
cloud of white smoke. At the FlG - 3 

same time the level of the water 

inside the bell-jar first falls a little and later rises; but while 
there is still a large volume of air left above the water on which 
the cork floats, the flame dies out and the burning ceases. By 

6 



_E 



17773^ 



The Burning of Substances — Oxygen 7 

the time the bell-jar and its content have become cold, the 
cloud has disappeared and the water has risen on the inside so 
that the volume of the remaining air is seen to be about four- 
fifths of the original volume. It follows that about one-fifth 
by volume of the air has disappeared. 

Further examination also shows that much of the phos- 
phorus still remains unburned. Why, then, should the burn- 
ing stop while there is still four-fifths by volume of the air left 
in the jar? The answer to this question may be made when 
we find that, try as we may, we cannot make phosphorus or 
anything else burn in the air remaining in the jar. We therefore 
conclude that the remaining air is different from common air. 
The correctness of this conclusion is supported by the fact that 
small animals, such as mice, suffocate at once if allowed to breath 
this remaining portion of the air. The facts just considered make 
it seem probable that one-fifth of the air is different from the 
balance, and that it is this portion which takes part in the burning 
of substances and which is necessary for the respiration of 
animals. 

Everyday experience would seem to indicate that wood, coal, 
paper, gasoline, etc., are completely destroyed when they are 
burned. Wood and coal leave a small amount of ash when 
burned, but nothing visible remains in the case of gasoline and 
other oils. Since we have found that water in the form of steam 
is invisible, it is possible that the substance burned may have 
passed into an invisible form and thus escaped notice. 

There are many substances which burn very readily and in 
so doing leave behind large amounts of ash; the experimental 
study of the burning of such substances leads to important con- 
clusions. We may now consider two typical cases of this sort. 

11. The Burning of Magnesium. — The metal magnesium, 
which is used in photographic flash lights, will burn very readily 
in air, either in the form of powder or thin ribbon. In either case 
we notice that a white ash is left. If we collect and weigh the ash 
from the burning of a weighed piece of magnesium ribbon we 
find that the ash weighs more than the original metal ribbon. 
The actual experiment is best carried out by placing about one 



8 



Introduction to General Chemistry 



O 



gram of magnesium, in the form of wire (Fig. 4) or ribbon, in a 
porcelain crucible, having a cover, and then weighing crucible and 

contents. The magnesium is then 
ignited and the cover so adjusted 
that some air can enter, but that 
the dense cloud of white smoke is 
largely held back in the crucible. 
After the burning is finished and 
the crucible has cooled and the 
whole is again weighed, it will be 
found that there has been a consider- 
able increase in weight. 
12. The Burning of Iron. — Iron powder or filings burn readily 
when thrown into a flame, and in a similar manner we find that 
the burned iron or iron ash, as we might possibly call it, is heavier 
than the original metal. In order to show this by experiment, 
we may suspend on one side of a balance (Fig. 5) a horseshoe 



Fig. 4 




Fig. 5 



magnet which has been dipped in iron filings, and counterpoise 
the magnet and adhering iron by adding small shot or sand to 
the other pan of the balance. By the application of a flame, the 
iron, which now presents a large surface to the air, may be 
ignited. As it burns with a dull' glow we observe a gradual 
increase in its weight, and, while there is no noticeable change 
in its volume, the cold residue, which we may call iron ash, is 



The Burning of Substances — Oxygen 



seen to have lost its metallic luster and taken on a dead black 
color. We find, thus, that iron ash is heavier than the iron burned. 
If we seek the cause of this increase in weight, we may get a 
hint when we remember that for the burning of a candle air is 
required, and that, moreover, part of the air disappeared when 
phosphorus was burned in it. What, then, becomes of the 
weight of the one-fifth of the air that disappeared? Is it added 
to the weight of the iron, so as to increase the weight of its ash? 
The facts presented in the next paragraph will furnish the 
required answers. 

13. Lavoisier's Experiment with Mercury. — An experiment 
which turned out to be one of the most important made in the 
early development of the 
science of chemistry was 
carried out by the great 
French chemist, Lavoisier, 
in the latter part of the 
eighteenth century. The 
arrangement in this classic 
experiment is shown in 
Fig. 6. The retort (the 
glass vessel with the long 
bent neck) was partly filled 
with mercury (quicksil- 
ver) ; the space above the mercury contained ordinary air, which 
also filled the bell-jar with which the neck of the retort communi- 
cated. The bell- jar stood in a shallow vessel containing mercury, 
which served to prevent outside air from passing into or out of 
the jar. The mercury in the retort was now heated by means of 
a charcoal stove for a period of several days. The heating first 
caused an expansion of the air; but as time went on a gradual 
contraction occurred, which entirely ceased after several days, 
whereupon the heating was stopped. The volume of the air left 
in the entire apparatus when brought to its original temperature 
and pressure was practically four -fifths of what it had been at the 
start. The surface of the mercury in the retort was found 
to be covered with a red powder, which may be considered 




Fig. 6 



io Introduction to General Chemistry 

as analogous to the white ash formed in the burning of 
magnesium or the black ash formed by the burning of iron 
filings. 

14. Heating the Red Ash of Mercury. — If we take some of the 
red ash of mercury, place it in a glass test tube, and heat it very 
strongly (Fig. 7), we find that it changes in a remarkable way: 
first it turns black, and then at red heat it gradually grows 
smaller, until after a few minutes none of it remains. At the 
same time, however, on the cooler part of the 
wall of the tube a silvery-looking coating has 
appeared, which when the tube has cooled 
may be brushed to the bottom of the tube, 
and is then readily seen to consist of drops of 
liquid mercury. Thus by heating the red 
powder to a higher temperature than that used 
in its formation, mercury is reproduced. But 
this is only half the story. 

The more important part remains to be 
told. Lavoisier reasoned about the matter 
somewhat as follows: If burning substances 
require air; if a part of the air disappears (in 
some cases at least) during burning; if in the 
burning of metals like magnesium and iron the ash is heavier than 
the metal burned; if, as is indeed a fact, air has weight; is it not 
possible that the burning substance unites with a part of the air 
to form a new kind of substance, and that this new substance, 
for example, magnesium ash, is heavier than the substance 
burned because it contains not only the latter but also a part of 
the air? Perhaps also the red ash formed by the gentle heating 
of mercury in contact with air is also made up of mercury and 
something taken from the air. Perhaps the one-fifth of the air 
that vanished has combined with the mercury to form the red 
ash. If all these suppositions are true, perhaps when the red 
ash was changed again into mercury by being strongly heated 
there was set free at the same time the part of the air which by 
originally uniting with the mercury produced the red ash. If 
all this were true, how could it be proved? Let us see. 




The Burning of Substances — Oxygen 



ii 



15. The Active Part of the Air: Oxygen.— The part of the air 
which disappeared may be just that part which causes substances 
to burn. If it were to be obtained pure, free from the inert four- 
fifths which does not support burning (combustion), it ought to 
support combustion far better than common air. This is a matter 
easily put to the test of experiment. Let us again heat some 
of the red powder in a test tube and at the same time thrust into 
the tube a burning wood splint. We see that it burns much more 
fiercely and brightly than in common air. Furthermore, if we 
have no flame, but only a tiny spark on the end of the splint, 
we see that when thrust into the tube above the heated red ash 
the spark bursts into a vigorous flame. The suppositions seem 
to be true. Lavoisier was led in this way to the discovery of 
the secret of the nature of burning. He called the gas formed 
by the heating of the red powder oxygen. This gas forms one- 
fifth by volume of the air and is the part of the air which is necessary 
for the burning of substances. The other four-fifths by volume 
of the air is inert; it does not support combustion; neither does 
it support the respiration of animals. Lavoisier called it azote; 
we call it nitrogen. 

16. The Properties of Oxygen. — Oxygen is an invisible gas 
like air; it has no odor and it supports combustion far better than 

does air. By the same method as that employed 
in the case of air (chap, i), we may readily find 
[jo that 1 liter of oxygen at a temperature of o° and 

76 cm. pressure weighs 1.43 g. It has, there- 
fore, a somewhat greater density than air, of 
which 1 liter weighs 1 . 29 g. Further evi- 
dence that the explanation of the nature of 
burning, given in the preceding paragraph, is the 
correct one is furnished by experiments which 
we may now consider. 

17. Burning Iron in Oxygen. — If we place a 
gram or two of iron filings and a minute piece of 
phosphorus on a piece of asbestos paper in the 
bottom of a 300-c.c. round-bottomed flask rilled with pure 
oxygen and fitted with a rubber stopper and a glass stopcock 




Fig. 8 



12 Introduction to General Chemistry 

(Fig. 8), we shall find that the weight of the flask with its con- 
tents does not change if by heating we cause the iron to burn 
in the oxygen. Now, we know that when iron burns, the prod- 
uct weighs more than the original iron. We know also that 
oxygen has weight and that the total weight of the flask with its 
contents has not changed during the burning. What then is the 
cause of the increase of the weight of iron when burned? If we 
open the stopcock while its open end is held under water, we find 
that the water nearly fills the flask. We must conclude that the 
oxygen has disappeared. Is it not reasonable to suppose that 
the ash resulting from the burning of iron is composed of the iron 
originally taken and the oxygen which has disappeared? Our 
experiment has shown that the weight of this ash is precisely the 
same as the combined weights of the iron and the oxygen which 
disappeared in the burning. 

It will readily be seen that the experiment with iron is similar 
to that made by Lavoisier with mercury — with the difference 
that iron burns rapidly, whereas mercury changes but slowly in 
oxygen. Furthermore, the fact that the red ash of mercury 
when strongly heated gives again mercury and oxygen makes 
it practically certain that the red ash was formed by the combina- 
tion or union of mercury with oxygen which composed part of 
the original air used in Lavoisier's experiment. Instead of iron, 
in the experiment described, we might have substituted magne- 
sium or phosphorus, or indeed any one of a large number of other 
substances. In each case the result would have been similar 
to that in the case of iron and oxygen and a similar conclusion 
would have been forced upon us. In all such cases we would 
conclude that the process of burning consists in the combination 
or union of gaseous oxygen with the solid substance burned to form 
the product of the combustion. 

18. Burning Charcoal in Oxygen. — If we put a piece of burn- 
ing charcoal into a bottle containing oxygen we notice that it 
burns even more rapidly in oxygen than in the air. In this case 
there is but a trifling amount of ash left compared with the 
amount of charcoal burned. In order to see whether an invisible 
product may have been produced we may make the following 
experiment. If we pour a little limewater into a bottle contain- 



The Burning of Substances — Oxygen 



13 



ing oxygen and shake the limewater with the oxygen we notice 
no change. If now we pour limewater into a bottle in which 
charcoal has been burned in oxygen and again shake the con- 
tainer, the limewater becomes milky in appearance. We must con- 
clude that some invisible substance, different from oxygen, has 
been produced in the latter case. If the burning of charcoal is 
thought to be analogous in nature to the burning of iron, then 
we might expect that the product would be something composed 
of carbon and oxygen and that its weight should be equal to the 
combined weights of the carbon burned and the oxygen taken up. 
We can get some evidence that this is the case by means of the 
following experiment. 

19. Carbon Dioxide. — A small quantity of charcoal is placed 
near one end of a hard glass tube, the other end of which contains 
pieces of caustic soda (Fig. 9). If we now weigh the tube, which 



-HIE 




Fig. 9 

may be fitted at the end containing the charcoal with a stopper 
and a small glass tube, and then cause the charcoal to burn in a 
stream of oxygen gas which we may pass through the tube, we 
shall find that there is an increase of weight, due to the fact that 
the product formed by burning the charcoal has been absorbed 
by the caustic soda in the tube. If we place some caustic soda 
in a beaker, dissolve it in water, and add some hydrochloric acid, 
we can see no marked change. If we treat the material from the 
charcoal experiment in the same way, we notice that a gas is 
given off when we pour the acid into the solution. A test of this 
gas with limewater shows that it behaves like that obtained 



14 



Introduction to General Chemistry 



when charcoal is burned directly in oxygen. The results of 
these experiments lead us to conclude that when charcoal is 
burned an invisible gas is produced, and that this gas is heavier 
than the charcoa burned; and, in fact, if charcoal had been 
burned in a closed vessel with oxygen, we should find that the 
weight of vessel and contents had not changed during the 
burning, and would be forced to conclude that the weight of 
the invisible product was just equal to the sum of the weights of 
the charcoal burned and the oxygen which had united with it. 
This gaseous product of the burning of charcoal was formerly 
called carbonic acid gas, but is now usually called carbon dioxide. 
20. Experiments with a Burning Candle. — We find by experi- 
ment that carbon dioxide is formed when wood, coal, illuminating 
gas, gasoline, etc., burn. We may easily show by the limewater 
test that it is also formed during the burning of a candle. We 

may also show that another well- 
known substance is produced when 
the candle burns. If we burn the 
candle under an inverted funnel 
connected by means of a glass 
tube with a U-tube which is cooled 
by immersion in a vessel of mercury 
and draw air through the funnel 
and U-tube we find that a colorless 
liquid .collects in the cold U-tube 
F IG IO (Fig. 10). This liquid is water. 

The burning of the candle gives, 
then, both carbon dioxide and water. We may readily show that 
the weight of the products of a burning candle, if these are suit- 
ably collected, is greater than the weight of the candle burned. 
To do this we make use of the arrangement shown in Fig. n. 
The candle is inclosed in a glass cylinder, closed below by a cork 
having threejor four holes for the admission of air. The top of 
the cylinder is filled with pieces of a solid substance (caustic 
potash) which readily absorbs both carbon dioxide and water, but 
not oxygen or nitrogen, the components of air. The apparatus 
thus arranged is suspended on one side of a balance and counter- 
poised. 




The Burning of Substances — Oxygen 



15 



The candle is now lighted and allowed to burn for ten or 
fifteen minutes, whereupon it will be found that the apparatus has 
become appreciably heavier. The increase in weight is due to the 
fact that the carbon dioxide and water formed weigh more than the 
candle burned. In fact, the excess weight is exactly the weight 
of the oxygen which has been consumed in the burning. Under 
ordinary circumstances the carbon dioxide and water escape 
our notice because both, the latter being in the form of steam, 
are invisible gases. 




Fig. 11 



21. The Law of the Conservation of Matter. — By the study 
of such facts as those discussed in the preceding paragraph and 
many others of a similar nature, Lavoisier arrived at the con- 
clusion that when a substance burns it unites with the oxygen of the 
air, and that the weight of the product is always exactly equal to the 
weight of the substance burned plus the weight of the oxygen which 
unites with the burned substance during the combustion. The 
product may be a solid, a liquid, or a gas. If it is a volatile 
liquid or a gas it usually escapes notice because it is invisible. 
Burning, therefore, consists in a union of the substance burned with 
oxygen. In this sense a substance which is burned is not destroyed; 
the material or matter composing it merely passes into another 
form, the quantity of matter in all cases being measured by its 
weight. These facts are briefly summed up in the statement that 
matter is indestructible, a statement which is frequently referred 
to as the Law of the Conservation of Matter. 



CHAPTER III 

PURE SUBSTANCES— ELEMENTS 

22. Bodies and Substanc.es. — We use the words "substance" 
and "body" in chemistry in very definite senses. We speak of 
things like watches or knives as " bodies." We say that the blade 
of the knife is steel, the handle is pearl. We say that a watch has 
a case of gold and a watch crystal of glass. We call steel, pearl, 
gold, and glass "substances." A substance is thus a particular 
kind of material, a body is an object which may be composed of 
one or many kinds of substances. Water, salt, and sugar are 
further examples of substances in the sense of this definition. 




Fig. 



23. Pure Substances. — We find that natural waters, as those 
of lakes, rivers, and springs, are not all alike. It now becomes 
important to discover the cause of the differences between 
waters from different sources. If we boil a quantity of lake 
water we find when the water has entirely disappeared that a 
solid residue is left. If the steam from the boiling water is con- 
densed by cooling it, as by means of a condenser (Fig. 12) through 
the outer tube of which a stream of cold water flows, we obtain 
what is called distilled water. If we now evaporate to dryness 
a quantity of this distilled water we find that no residue is left. 
If we prepare distilled water from any natural water we find that 

16 



Pure Substances — Elements 



17 



it will always evaporate completely, leaving no solid residue. 
We find further that different kinds of natural water leave differ- 
ent proportions of solid residue upon evaporation and that the 
nature of the solid material left also differs in different cases, but 
that the distilled water in one case cannot be distinguished in 
any way from that obtained in another. We say then that 
distilled water is pure water, a pure substance, and the natural 
waters are not pure water, but that they contain dissolved foreign 
substances. If the natural water is muddy, that is, if it is not 

clear, the foreign material 
which causes it to appear 
muddy can be separated 
by filtration (Fig. 13), a 
process in which the liquid 
is allowed to seep through 
a piece of filter paper 
folded so as to fit snugly 
into a funnel. The mud 
remains on the filter paper. 
However, filtration will not 




H 



D 



Fig. 13 



remove any of the dissolved material, but only that which is 
suspended in the water. 

24. Pure Salt Made from Rock Salt. — Common salt is found 
in nature as a mineral known as rock salt. We find that different 
samples of rock salt differ in color, taste, specific gravity, and in 
other ways. If we mix rock salt with water we find that a large 
part dissolves in the water. In general a small amount of 
material, sand, etc., will not dissolve, even though we take a large 
amount of water. If we filter the solution we separate the water 
and dissolved material from the part which has not dissolved. 
That which runs through the filter paper is called the filtrate ; it 
is the solution of the salt in water. If we boil away the water 
we find that the salt is left in the solid form and that the material 
is now free from color, that is, that it is white, and that it will 
dissolve completely in water. The salt so prepared is purer than 
the rock salt taken. Just as it is possible to prepare pure water 
from any natural water, so analogously it is possible to prepare 



Introduction to General Chemistry 



pure salt from any natural salt. Pure salt is always exactly the 
same in taste, color, specific gravity, etc., from whatever source 
it may have been obtained. The process for the purification 
of salt, described in the statement above, gives in all cases a 
much purer product than the original rock salt — pure enough 
for table use, but not a perfectly pure substance. It still con- 
tains very small amounts of some foreign substances; but even 
these can be removed by well-known methods which the student 
will learn later. A pure substance is a substance which consists 
of one sort of material. It always has definite physical proper- 
ties, from whatever source it may be obtained. 

25. Decomposition of Substances. — It was found that the red 
ash formed by heating mercury in contact with air was changed, 
upon being heated still more, into mercury and oxygen. We say 
in this case that the red ash of mercury has been decomposed into 

mercury and oxygen. We can 
accomplish the decomposition of 
many substances in an equally 
simple fashion. 

We will now consider a few 
such cases as illustrations. 

26. Decomposition of Sal 
Soda. — If we place in a test tube 
a crystal of common washing-soda, 
also known as sal soda, and heat it 
gently over a Bunsen flame, we find 
that water is produced as steam, 
and that it condenses in the cold 
end of the test tube. An opaque 
solid is left in place of the clear 
crystal of sal soda taken. We say 
that the sal soda has been decom- 
posed into dry soda and water. It 
would be easy to show that the weight of the water and dry soda 
formed is equal to the weight of sal soda taken. In other 
words, the sal soda has been decomposed into dry soda and 
water. 




Pure Substances — Elements 19 

27. Electrolysis of Water. — If we pass an electric current 
through some water (Fig. 14) to which we have added a few drops 
of sulfuric acid, we find that gases are produced at the platinum 
electrodes. The decomposition of a substance by an electric 
current is called electrolysis. If we collect each of these gases 
separately we find that one of them is oxygen. The other gas, 
the volume of which is double that of the oxygen, has quite differ- 
ent properties; it is called hydrogen. If we bring a lighted 
splinter into the oxygen, the splinter continues to burn with 
increased brilliancy and rapidity. If we repeat this test with 
hydrogen, we find that the hydrogen itself catches fire, just as 
illuminating gas would do, and that the splinter itself no longer 
burns in the hydrogen gas. These facts may be concisely stated 
by saying that oxygen supports combustion, while hydrogen burns 
but does not support combustion. It would be possible to show 
by experiment that the weight of the water decreases during the 
passage of the electric current through it, and that this decrease 
in weight is just equal to the combined weights of the oxygen and 
hydrogen formed. The total amount of sulfuric acid added to 
the water remains in the water at the end of the electrolysis and 
would serve to promote the decomposition of any desired amount 
of water. The complete explanation of the behavior- of the 
sulfuric acid cannot be given at this point, but we know that the 
hydrogen and oxygen formed come exclusively from the water 
and not from the acid nor the platinum nor the glass of the 
vessel used. We conclude that water is decomposed by the electric 
current into hydrogen and oxygen. Therefore we may say that 
water is composed of hydrogen and oxygen or that water is a 
compound of hydrogen and oxygen. As a matter of fact, when 
hydrogen burns in air water is formed. If a cold beaker is held 
over a jet of burning hydrogen, water will be seen to condense in 
a mist on the surface of the beaker. 

28. Magnesium Burned in Steam. — That water is composed 
of oxygen and hydrogen may be shown in many other ways, one 
of which is the following. When a piece of magnesium ribbon 
burns in air the magnesium unites with the oxygen of the air 
to form a white solid which we call magnesium oxide. Xow, 



20 



Introduction to General Chemistry 



magnesium will also burn in steam (Fig. 15) nearly as readily 
as it does in the air or in pure oxygen, and we find that the 
white solid which is again formed is also magnesium oxide. 
In addition, hydrogen gas is produced and may easily be col- 
lected over water. Since magnesium oxide is composed of 
magnesium and oxygen, and we obtain from magnesium and 
water magnesium oxide and hydrogen, we are again led to the 
conclusion that water is composed of hydrogen and oxygen. 

29. Steam Passed over Hot Iron. — An entirely analogous 
experiment may be carried out with iron and steam. In this 




Fig. 15 

case iron turnings or fine iron wire is strongly heated in an iron or 
glass tube (Fig. 16). When steam is passed through the tube, 
iron oxide and hydrogen are produced, a result which leads to the 
same conclusion as before regarding the composition of water. 

30. Magnesium Burned in Carbon Dioxide. — The composi- 
tion of carbon dioxide may be discovered by burning magnesium 
in this gas. We find that magnesium oxide and a product 
resembling charcoal are formed. The latter substance is carbon, 
of which charcoal is a nearly pure form. We conclude, therefore, 
that carbon dioxide is composed of carbon and oxygen or is a 
compound of carbon and oxygen. 

The facts already considered lead to the conclusion that the 
red ash obtained when mercury is heated gently in air is com- 



Pure Substances — Elements 



21 



posed of mercury and oxygen; briefly, that it is a compound of 
mercury and oxygen — a fact represented by the chemical name 
of the red ash, mercuric oxide. 

31. Elements. — The substances mercury, oxygen, hydrogen, 
and carbon have never been decomposed into simpler substances. 
We say that hydrogen and oxygen are the elements of which 
water is composed; that carbon and oxygen are the elements com- 
posing carbon dioxide. 

We may discover of what elements a substance is composed in 
two ways: either by the decomposition of the substance into the 




Fig. 16 

simpler ones which compose it — the process called analysis, or 
by causing known elements to unite in the formation of the 
original — the process called synthesis. As a result of the 
electrolysis of water we have concluded that water is composed 
of hydrogen and oxygen. This conclusion may now be tested 
by seeing whether water can be obtained from hydrogen and 
oxygen. We found that hydrogen burns readily. If we burn 
a jet of hydrogen under an inverted funnel and draw the product 
through a cooled U-tube, as in the experiment with the candle, 
we shall find that liquid water collects in the U-tube and that 
the most careful search fails to reveal any other substance as 
the product of the burning of hydrogen in air or in pure oxygen. 
Water is, therefore, a compound of the elements hydrogen and 



22 Introduction to General Chemistry 

oxygen. Since the burning of charcoal, which is a nearly pure 
form of the element carbon, gives carbon dioxide and nothing 
else, we know that carbon dioxide is a compound of the elements 
carbon and oxygen. 

32. The Burning of Copper; Copper Oxide. — If the metal 
copper, in the form of fine wire or filings, is heated in air or in 
oxygen, it is slowly changed into a black substance quite different 
in appearance from metallic copper; but during this change we 
do not observe the production of any light. By means of the 
balance we may find that the black substance formed is heavier 
than the copper taken, and we at once suspect that the copper 
has united with oxygen to form a compound. If the heating 
of the copper had been carried out in a sealed glass vessel con- 
taining oxygen, as in the earlier experiment with iron powder, 
it would have been found that gaseous oxygen had disappeared 
and that the weight of the black product was exactly equal to the 
weight of the copper taken plus the weight of the gaseous oxygen 
which had disappeared. The black substance would seem, there- 
fore, to be a compound of copper and oxygen. We know that 
when the red mercury oxide is strongly heated it Is decomposed 
into mercury and oxygen. If we heat the black product from 
copper to the highest temperature we can obtain with the Bunsen 
burner, we find that it remains unaltered in weight and appear- 
ance and that no oxygen is given off. This fact might lead us 
to suspect that the black substance is not a compound of copper 
and oxygen, since its behavior is not analogous to that of mercury 
oxide. In this connection the following experiment will prove 
of interest. 

33. Hydrogen Passed over Hot Copper Oxide.— If we put two 
or three grams of the black copper product in a porcelain boat in 
a "hard" or difficultly fusible glass tube, heat the tube and con- 
tents by means of a Bunsen flame, and then pass a current of 
hydrogen through the tube, we observe that the solid glows or 
seems to burn (Fig. 17). At the same time we notice that liquid 
water condenses in the colder part of the glass tube. After a few 
minutes the glow has disappeared, even though the stream of 
hydrogen has continued. At this point the heating may be 



Pure Substances — Elements 23 

discontinued and the solid which is left in the boat allowed to 
cool in the stream of hydrogen gas. We now observe that the 
solid has the appearance and properties of metallic copper, which 
in fact it is. However, the copper is not in a single compact 
lump, for a reason which must be evident. Metallic copper can 
be melted, but the melting-point is a very much higher tempera- 
ture than that attained in the preceding experiment. Only by 
heating the copper to a point above this melting temperature 
could the material be obtained in a single lump. This could 
easily be accomplished by directing an intense blowpipe flame 
upon the metal particles contained in the porcelain boat. 

We may now consider the nature of the changes which occured 
in this experiment. Since we obtained water and copper, and 



^SM 



r? 



Fig. 17 

since we know that water is a compound of the element hydrogen 
with oxygen, we conclude that the oxygen was originally united 
with the copper and that the black substance must have been a 
compound of copper and oxygen. This substance is called 
copper oxide. We might express the result in the following 
simple fashion: Copper Oxide + Hydrogen -> Wat er+ Copper; or 
instead of " Water" we might write "Hydrogen Oxide," the true 
chemical name for Water. This statement would then show at 
a glance the nature of the chemical change which had occurred. 
34. Discovery of the Elements Composing a Substance : the 
Analysis of Malachite. — There is an almost innumerable variety 
of bodies on and in the earth, but these are composed of a very 
much smaller number of definite chemical substances. However, 
the number of definite substances is still very great, many 
thousands having been carefully described. Chemistry has for 
its object the systematic study of pure substances, their properties. 



24 Introduction to General Chemistry 

and their behavior toward one another. Happily the study of this 
immense number of substances is greatly simplified by the fact 
that they are all made up of a relatively small number of elements. 
The way in which the elements composing a substance of un- 
known composition are discovered may be illustrated by means 
of an experiment with the mineral known as malachite. Mala- 
chite is a beautiful crystalline substance often used as an orna- 
mental stone and also as one of the sources from which a familiar 
metal is obtained. If we place in a test tube, fitted with a cork 
and a bent glass tube, a few grams of malachite and heat the 
substance gently in a flame, we notice that a change in color 
from green to black occurs and at the same time that water 
condenses in the colder part of the glass tube and a gas is also 
given off. If we pass this gas into limewater we find that it 
behaves like carbon dioxide, which in fact it is. By means of 
the balance we might find that the combined weights of the 
carbon dioxide, water, and black product equal the weight of 
the original malachite. Since we know of what elements carbon 
dioxide 'and water are composed, it only remains to find the 
composition of the black substance in order to have a complete 
knowledge of the elements composing malachite. If this black 
substance were heated in a stream of hydrogen, it would be found 
to yield water and a red metallic-looking substance which could 
easily be recognized as copper. Therefore, the black substance 
must have been copper oxide. The results may then readily 
be interpreted. Malachite when heated is decomposed into 
carbon dioxide, water, and copper oxide. Knowing as we do the 
elements composing each of these three products, we are led to 
the conclusion that malachite is a compound of the elements 
carbon, oxygen, hydrogen, and copper. Chemists have so far been 
unable to decompose copper into anything simpler. It is, 
therefore, known as an elementary or simple substance, and we 
say that malachite is a compound of the four elements, carbon, 
hydrogen, oxygen, and copper. 

35. Some Common Elements. — The total number of known 
elements is about eighty-five, of which less than thirty are com- 
mon. In the following partial list of commoner elements, the 



Pure Substances — Elements 25 

student will find the names of ten or twelve familiar metals. 
Carbon and sulfur, which are well known to everyone, are not 
metals; they are classed as non-metals. 

A FEW COMMON ELEMENTS 

Silver Copper Nickel Carbon 

Gold Lead Magnesium Sulfur 

Platinum Tin Zinc • Oxygen 

Iron Aluminum Mercury Hydrogen 



CHAPTER IV 
THE LAW OF DEFINITE COMPOSITION 

36. The Percentage Composition of Water. — We have 
already seen that when an electric current was passed through 
water, the latter was decomposed into two gases, hydrogen and 
oxygen. It was found that the volume of the hydrogen was 
double that of the oxygen obtained in the electrolysis. This was 
not a matter of accident, for it is always found that the same 
result is obtained whenever water is electrolyzed. Since water 
is composed only of hydrogen and oxygen, we may calculate the 
percentages of hydrogen and oxygen by weight if we know the 
weight of a liter of each of these gases. Direct weighing of the 
gases has shown that 1 liter of hydrogen weighs 0.090 g. and 
1 liter of oxygen weighs 1.429 g., the gases being weighed at o° 
and 76 cm. pressure. From these figures it is easy to calculate 
that water is composed of 1 1 . 2 per cent of hydrogen and 88 . 8 
per cent of oxygen by weight. Pure water prepared from any 
source whatever always has exactly this composition. 

The percentage composition of water may also be found in 
another way. It was found in section 33 that water and copper 
are formed when hydrogen is passed over heated copper oxide. 
If this experiment be carried out with a weighed quantity of 
copper oxide, and the weight of copper which remains after the 
experiment is found, the difference in the two weights will repre- 
sent the weight of oxygen contained in the water which has been 
formed. If the weight of the water is determined, then the 
percentage of oxygen in water may readily be calculated. In 
this case we find precisely the same result as that given in the 
preceding paragraph. 

The details of the experiment are as follows. 

37. The Quantitative Synthesis of Water. — About one gram 
of pure copper oxide is placed in a weighed porcelain boat and 
heated sufficiently to drive off the moisture which it may con- 

26 



The Law of Definite Composition 



27 



tain. 1 The boat and contents are weighed as soon as cool and 
placed at once in a hard glass tube. This tube (Fig. 18) is con- 
nected at each end with U -tubes rilled with calcium chloride, a 
substance that absorbs water with great readiness. One of 
these U -tubes is connected with a source of hydrogen gas and 
serves to remove all moisture (water vapor) from the hydrogen. 
The other U-tube will serve to absorb the water formed in the 
chemical reaction between the copper oxide and the hydrogen. 




5 5 5 



:m. 




Fig. 18 



This second U-tube is accurately weighed at the beginning of 
the experiment. 

When all is ready, the stream of hydrogen is started and con- 
tinued until all air is driven from the tubes. The tube contain- 
ing the boat is now heated until the reaction begins, and kept hot 
enough beyond the boat to prevent the condensation of the 
steam formed, which is carried by the stream of hydrogen into 
the weighed U-tube. 

When all the copper oxide has been changed into copper and 
the water has all been driven over into the U-tube, the heating 
is discontinued and the copper allowed to cool in a stream of 
hydrogen. The hydrogen is then driven out by a stream of air, 
and the U-tube detached and weighed. The object in repla- 
cing the hydrogen by air is readily understood when one recalls 
that hydrogen is far lighter than air. Therefore the weight of 
the tube filled with hydrogen would be appreciably less than 
if it is rilled with air. The increase in weight is the weight 
of the water formed. The boat containing the copper is also 
weighed. The loss in weight is the weight of oxygen contained 

1 Most substances, especially if porous or in the form of powder, absorb more 
or less moisture from the air. 



28 Introduction to General Chemistry 

in the water formed. The results of an actual experiment 

were as follows: 

Boat and copper oxide 9 . 523 g. 

Boat 8.451 

Copper oxide 1 . 072 g. 

Boat and copper 9 .311 g. 

Boat 8.451 

Copper o . 860 g. 

Tube and water 18 . 665 g. 

Tube 18.426 

Water o . 239 g. 

Since 1.072 g.— 0.860 g. =0.212 g., we conclude that o.239g. 
of water was formed from o. 212 g. of oxygen, which at the begin- 
ning was in combination with the copper in the form of copper 
oxide. Therefore water consists of 0.212 g.-i-o. 239 g. = 0.887 
= 88 . 7 per cent oxygen. The difference between the weight 
of water formed and that of the oxygen' used is the weight of 
hydrogen, which is 0.239 g. — 0.212 g.= 0.02 7 g. This is readily 
found to be 11. 3 per cent of the weight of the water. Very 
carefully peformed experiments, made in this way, show that 
water contains 88 . 8 per cent by weight of oxygen and 1 1 . 2 per 
cent of hydrogen; the difference of 0.1 per cent between the 
values found in the lecture experiment quoted and those ob- 
tained in the most accurate experiments made by skilled chemists 
working with greatest care and under ideal conditions is due to 
the experimental errors in the rather crude lecture experiment. 

38. The Percentage Composition of Copper Oxide. — It is also 
easy to see that we may find the percentage composition of copper 
oxide from the data just considered. Thus 1.072 g. of copper 
oxide gave o. 860 g. of copper by loss of o. 212 g. of oxygen; from 
which we find that copper oxide is composed of 80 . 2 per cent 
copper and 19.8 per cent oxygen. The most accurate experi- 
ments made in this way give 79.9 per cent copper and 20.1 
per cent oxygen, the difference being due to experimental error 
in the lecture experiment. Pure copper oxide always has exactly 
the composition shown by these figures. 



The Law of Definite Composition 29 

39. The Percentage Composition of Carbon Dioxide. — We 

have found that carbon in the form of charcoal burns readily in 
air or in oxygen with the formation of a colorless gas called carbon 
dioxide. The percentage composition of carbon dioxide may 
be found by burning a known weight of pure carbon in oxygen 
gas and finding the weight of carbon dioxide formed. It will be 
recalled that carbon dioxide is easily absorbed by solid caustic 
soda. It is also readily absorbed by a solution of caustic potash 
in water, while neither oxygen nor air is absorbed by such a solu- 
tion. If the gases formed by the burning of carbon in a stream 
of oxygen are passed through a suitable bulb containing caustic 
potash solution, all of the carbon dioxide will be retained by the 
solution and the oxygen will pass through unabsorbed. The 
increase in weight of the bulb will represent the weight of the 
carbon dioxide formed by the burning of the carbon. 




Fig. 19 

The arrangement of the apparatus is shown in Fig. 19. 
About o. 2 g. of pure carbon, made from sugar, is contained in a 
porcelain boat which is placed in a hard glass tube connected at 
one end with a supply of pure oxygen and at the other with a 
calcium chloride tube and a weighed potash bulb, which contains 
a 30 per cent solution of caustic potash. The middle part of the 
tube should contain a column of copper oxide, to insure the 
complete conversion of the carbon into carbon dioxide. The 
calcium chloride tube serves to catch any moisture present. 
The carbon is ignited by heating the tube with a gas burner; 
after the carbon has completely burned and all of the carbon 
dioxide formed has been driven over into the potash bulb by 
the stream of oxygen, a slow stream of air is blown or drawn 
through the apparatus to replace the oxygen by air. The 
potash bulb is then detached and weighed. In an actual lecture 
experiment o. 194 g. of carbon yielded o. 701 g. of carbon dioxide; 



30 Introduction to General Chemistry 

from which we find that this gas contains 27.6 per cent of carbon 
and 72.4 per cent of oxygen. The most accurate experiments 
of skilled chemists show the correct percentages to be 27.3 per 
cent carbon and 72.7 per cent oxygen. 

40. The Action of Sodium on Water : Caustic Soda. — It is a 
matter of importance to know the exact percentage composition 
of pure substances and a great variety of methods must be 
employed in the making of such determinations. It often 
happens that the method which would seem to be most direct 
and desirable is not practicable because the violence of the inter- 
action of the elements which we bring together would cause loss 
of some of the material taken. This may be illustrated by an 
experiment with the element sodium. If we throw a piece of 
this metal upon water, we observe that the action is a violent 
one which ordinarily ends in an explosion that throws part, of 
the substance out of the beaker in which it was contained. We 
may carry out the same reaction without loss of material and 
obtain precisely the same product if the piece of sodium is 
exposed to water vapor instead of being thrown upon liquid 
water. In this case the reaction requires much more time, but 
it proceeds quietly and without loss of material. The white 
solid so obtained is caustic soda. 

41. The Action of Hydrochloric Acid on Caustic Soda: Com- 
mon Salt. — If we add to a solution of caustic soda contained in a 
beaker a sufficient amount of pure hydrochloric acid and evap- 
orate the resulting solution to dryness, we find that the product 
is one with which we are well acquainted. It is nothing more 
nor less than common salt, and if the materials used are all pure 
the product will be chemically pure salt. We discover in this 
way that the metal sodium is one of the constituents of common 
salt. In fact, metallic sodium may be obtained by the elec- 
trolysis of molten salt, although this is not the most satisfactory 
method of making this metal. The percentage of sodium in salt 
may readily be found if the weights of sodium taken and of salt 
obtained are determined. 

42. The Percentage of Sodium in Common Salt. — In an 
actual experiment 0.483 g. of metallic sodium was weighed in a 



The Law of Definite Composition 31 

stoppered test tube (to prevent action of the moisture of the air) . 

The sodium was placed on a strip of silver foil which rested on the 

edges of a small porcelain dish containing about 10 c.c. of water, 

and covered with a beaker. In the course of a few hours the 

sodium had reacted completely with the water vapor to form a 

solution of caustic soda which dripped into the dish. A little 

of the solution adhering to the foil was rinsed into the dish with 

a little water. Sufficient pure hydrochloric acid was then added 

and the solution evaporated by steam heat in the manner shown 

in Fig. 20. The beaker contained ordinary water. By this 

mode of evaporation of the solution in the dish we avoid loss by 

spattering that would occur if we should 

boil the solution by heating the dish V^7 

directly with the flame. When the salt 

appeared to be dry, the dish was heated 

very cautiously with the direct flame, to 

drive off the small amount of remaining 

water. When cold, the dish and contents 

were weighed. It was found in this way — 

that 0.483 g. of sodium gave 1.217 g. of fig. 20 

common salt, which indicated that salt 

contains 39.7 per cent of sodium. The correct result is 39.4 

per cent. 

43. The Electrolysis of Hydrochloric Acid: Chlorine. — It is, 
of course, obvious that the sodium in common salt must be com- 
bined with one or more elements and the student will readily 
guess that a clue to the other constituents of common salt may 
be gained by a knowledge of the constituents of hydrochloric 
acid. If we pass an electric current through a concentrated solu- 
tion of hydrochloric acid contained in the apparatus shown in 
Fig. 21, we find that two gaseous products are obtained, the 
volumes of which are practically equal. One of these is colorless. 
It is lighter than air and burns with a hot but non-luminous 
flame and in so doing yields water; these properties show the 
colorless gas to be hydrogen. The other gas is pale yellow in 
color; it is heavier than air, one liter weighing 3. 22 g., and has 
an exceedingly disagreeable, irritating odor. This gas is known 



32 



Introduction to General Chemistry 



as chlorine. Inasmuch as chlorine has never been separated 
into simpler substances, we conclude that it is an element. 

44. The Union of Hydrogen and Chlorine: Hydrogen 
Chloride Gas. — Since the hydrochloric acid which was elec- 
trolyzed contained water, we should not be warranted in con- 
cluding that hydrogen is a constituent of hydrochloric acid; 

for, as we know, hydrogen is also one of 
the constituents of water. If we bring 
together equal volumes of the gases 
hydrogen and chlorine and allow them 
to mix, and if we allow the vessel to 
stand in diffused light for a day or 
two, we notice that the yellow color 
of the chlorine has disappeared. We 
find that a colorless gas remains which 
dissolves with the greatest ease in 
water, and that neither hydrogen nor 
chlorine is left. Since water which has 
dissolved the gas has all of the proper- 
ties of a solution of pure hydrochloric 
acid, we interpret the results as show- 
ing that equal volumes of hydrogen 
and chlorine gases combine to form 
a new gas which we call hydrogen 
chloride gas, and that the latter when 
dissolved in water constitutes hydro- 
chloric acid. Hydrogen chloride gas may be distinguished from 
the other gases which we have met in several ways, notably by 
its marked, choking odor, by the fact that it fumes or gives a 
white cloud in moist air, and it dissolves with great ease in 
water, as well as in several other ways. 

45. Salt a Compound of Sodium and Chlorine. — The fact that 
hydrochloric acid is known to be a compound of chlorine suggests 
that common salt may also contain this element. This is in 
fact the case. It can readily be shown by experiment that 
common salt results from the union of chlorine gas with metallic 
sodium. Inasmuch as nothing else is needed and no other 




Fig. 21; 
apparatus. 



The Law of Definite Composition 33 

product than salt is formed, we must conclude that salt is a 
compound of the elements sodium and chlorine. This fact is 
indicated by the chemical name of common salt, sodium 
chloride. Since salt contains 39.4 per cent of sodium, the 
percentage of chlorine must be 60.6. 

46. The Law of Definite Composition. — The preceding para- 
graphs of this chapter are intended to illustrate how we may 
arrive at a knowledge of the nature and percentage by weight 
of each element entering into the composition of a pure substance. 
It is possible, by well-known methods, to do this for all pure 
substances. As a result of countless thousands of such quantita- 
tive experiments made by chemists, the conclusion has been 
reached that the percentage composition of every pure substance 
is perfectly definite for that substance and is found to be the same 
by whatever method we may make the determination. This is 
one of the most important laws of chemistry. It is usually 
spoken of as the Law of Definite Composition or of Definite 
Proportions. This explains why a pure substance always has 
definite properties, from whatever source it may be obtained. 

47. Hydrogen and Its Gaseous Compounds. — We have 
already become acquainted with hydrogen and one of its gaseous 
compounds, hydrogen chloride, a water solution of which is 
known as hydrochloric acid. Hydrogen forms many compounds 
which are gaseous at ordinary temperatures. We shall now take 
up a study of some of these, with the object in view, first, of dis- 
covering the nature of the other element combined with the 
hydrogen; secondly, of discovering the percentage composition; 
and, finally, of disclosing a very remarkable relation between the 
weights of hydrogen contained in equal volumes of these gases. 

48. Hydrogen Chloride. — We have found that equal volumes 
of hydrogen and chlorine combined to form hydrogen chloride 
gas. Since we know that 1 liter of hydrogen weighs 0.090 g. 
and that 1 liter of chlorine weighs 3. 220 g., we find by calcula- 
tion that hydrogen chloride contains 2 . 76 per cent by weight of 
hydrogen. By direct weighing of pure hydrogen chloride gas 
it is found that 1 liter weighs 1.642 g. Since 2.76 per cent of 
1.642 g. is 0.045 g-> ^ follows that 1 liter of hydrogen chloride 



— ^^^^^ 



34 Introduction to General Chemistry 

gas contains 0.045 g. °f combined hydrogen. It has already 
been stated that 1 liter of hydrogen gas weighs 0.090 g., which 
weight we see is exactly double the weight of hydrogen in 1 liter of 
hydrogen chloride gas. 

49. Acetylene: a Compound of Carbon and Hydrogen. — 
Let us next consider the gas acetylene which is extensively used 
for illumination. This gas is obtained by allowing water to 
drop on calcium carbide. We find that it is a colorless gas with 
a peculiar odor. Everyone knows that it burns in air, giving an 
exceedingly bright flame. If we collect and test the products 
coming from the acetylene flame we find carbon dioxide and 
water. We find the same products and no others when acetylene 
is burned in pure oxygen gas, and therefore conclude that carbon 




Fig. 22 

and hydrogen are constituents of acetylene; but the experiment 
obviously does not decide whether oxygen is or is not also a 
constituent of acetylene. This question could be decided if we 
knew the percentages of carbon and hydrogen in the gas. 

50. The Analysis of Acetylene. — We may find the per- 
centages of carbon and hydrogen by means of the following 
experiment. A tube of hard glass a centimeter or more in 
diameter and 30 cm. long (Fig. 22) is partly filled with pure dry 
copper oxide. The tube is then heated red hot and a measured 
volume of acetylene at a known temperature and pressure is 
caused to pass through the tube and over the heated copper 
oxide. It is found that carbon dioxide and water are formed 
and that part of the copper oxide is changed into metallic copper. 
A U-tube filled with calcium chloride, for the absorption of the 
water formed, is attached to the exit of the hard glass tube. 
Beyond this, attached by rubber tubing, we have a bulb contain- 
ing caustic potash solution to absorb the carbon dioxide. After 



The Law of Definite Composition 35 

all of the acetylene has been driven over into the combustion 
tube holding the copper oxide, by allowing mercury from the 
attached reservoir slowly to displace the acetylene, a slow stream 
of pure dry oxygen is passed into the combustion tube to insure 
the complete burning of the carbon of the acetylene. Finally, 
the oxygen is displaced by a stream of air. 

The increase in weight of the calcium chloride tube represents 
the weight of water formed. Similarly the increase in weight 
of the caustic potash bulb represents the weight of carbon dioxide 
obtained. Now we. know that water contains 11. 2 per cent of 
hydrogen and that carbon dioxide contains 27.3 per cent of 
carbon. We may then calculate the weights of hydrogen and 
carbon corresponding to the weights of water and carbon dioxide 
obtained. If we know that 1 liter of acetylene under standard 
conditions, that is, at o° and 76 cm. P, weighs 1 . 190 g., we have 
all the data needed to enable us to calculate the percentages of 
hydrogen and carbon in acetylene. In an actual lecture experi- 
ment 200 c.c. of pure dry acetylene at 18 and 75.4 cm. gave 
o.isog. of water and 0.751 g. of carbon dioxide. From the 
data above we find that the weight of the acetylene taken was 
o. 222 g., and that the weights of hydrogen and carbon contained 
in the water and carbon dioxide respectively were 0.0168 g. and 
0.205 g-> respectively. Therefore acetylene contains (according 
to this analysis) 7.5 per cent of hydrogen and 92.3 per cent of 
carbon. The correct percentages are 7 . 7 and 92.3 respectively; 
and since the sum of these percentages is 100, we know that 
hydrogen and carbon are the only elements contained in acetylene. 
We may also calculate from the same data the weight of com- 
bined hydrogen in one liter of acetylene under standard condi- 
tions. We find in this way 0.090 g. of hydrogen. 

51. Ammonia. — Let us next take up the study of ammonia. 
Common household ammonia, which is familiar to everyone, is a 
solution in water of the substance, ammonia, which is a gas at 
ordinary temperature and pressure. If we warm such a solution 
of ammonia, a gas having an intense odor is given off. When 
this gas, ammonia, is strongly compressed, it condenses to a 
colorless liquid which we speak of as liquid ammonia. This is 



36 Introduction to General Chemistry 

a commercial article which is shipped in heavy steel cylinders six 
feet long and a foot in diameter. The liquid ammonia exists 
under considerable pressure in such cylinders. If the valve of 
the cylinder is opened gaseous ammonia escapes. We may use 
a small cylinder of liquid ammonia as a convenient source of 
ammonia gas. 

If we fill a glass cylinder with mercury, invert it in a dish of 
mercury, and allow ammonia gas to escape under the mouth of 
the cylinder, the mercury is displaced by the ammonia gas. We 
notice that the gas is invisible, like air. It is to be distinguished 
from air, however, by its intense odor, as well as in other ways. 
If we dip the mouth of the cylinder, which has been closed by a 
glass plate, into a vessel of water, we find that the water rushes 
into the cylinder almost as readily as if the space were a vacuum. 
An examination of the water now shows that it has new prop- 
erties. The water now has the odor of ammonia, it has a 
peculiar disagreeable taste, and changes the color of immersed 
red litmus paper blue. If we bring a burning candle into a 
cylinder of ammonia the flame of the candle is extinguished but 
the ammonia does not take fire. These properties distinguish 
ammonia from oxygen, hydrogen, and acetylene. 

52. Ammonia a Compound of Nitrogen and Hydrogen. — We 
may now inquire, What is tne chemical composition of ammonia? 
Is it an elementary substance or a compound, and, if a compound, 
of what elements is it composed? If ammonia gas is passed 
through a heated glass tube containing copper oxide we observe 
that a colorless liquid condenses in the cold part of the tube. 
This liquid proves to be water. We find also that a colorless, 
odorless gas is formed. If we pass this gas into limewater we 
observe no result and conclude, therefore, that this gas is not 
carbon dioxide. We find that the gas is not appreciably soluble 
in water, so that it cannot be unchanged ammonia gas. If we 
test the gas with a burning candle we find that it neither burns 
nor supports combustion. The student will doubtless recall 
(10) that this gas has just those properties which the portion 
of the air left after the removal of oxygen by mercury or phos- 
phorus possesses. It would seem, therefore, to be nitrogen. 



The Law of Definite Composition 37 

The identity of the gas with nitrogen is confirmed by a deter- 
mination of the density; whereupon it is found that a liter 
weighs i.25ig. Since water and copper were formed from 
ammonia and copper oxide, we conclude that ammonia has 
furnished the hydrogen which united with the oxygen supplied 
by the copper oxide to form the water obtained in the preceding 
experiment. Ammonia must be a compound containing nitrogen 
and hydrogen. It has been shown in many ways by experiments, 
which we need not consider at present, that nitrogen and hydro- 
gen are the only constituents of ammonia. 

53. The Percentage Composition of Ammonia. — The per- 
centage of hydrogen in ammonia may be found by carrying 
out the experiment above described with a known volume of 
ammonia measured at a known temperature and pressure. If 
we cause the ammonia to pass through the heated copper oxide 
tube, driving out water vapor completely by means of air after 
all of the ammonia has passed into the tube, and if the products 
are caused to pass through a calcium chloride tube connected 
to the copper oxide tube as in the determination of the composi- 
tion of acetylene, the increase in weight of the calcium chloride 
tube gives us the weight of water formed from the hydrogen of 
the ammonia used. Knowing as we do the percentage of hydro- 
gen in water, if we know the weight of a liter of ammonia gas 
(o.772g.) we may calculate the percentage of hydrogen in 
ammonia and also the weight of combined hydrogen in 1 liter of 
ammonia gas measured under standard conditions. We find 
this latter weight to be o. 135 g. 

54. Methane, Another Compound of Carbon and Hydro- 
gen. — The chief component of natural gas is a substance called 
methane. This same gas methane often escapes in bubbles when 
the decaying vegetable matter in marshes is disturbed. For this 
reason methane is also known as marsh gas. We may prepare 
methane artificially in the laboratory by methods which we need 
not now discuss. It may be collected over water, as its solu- 
bility 'in water is slight. We note that it is a colorless gas, that 
it is lighter than air, since the gas will escape rapidly from an 
open cylinder when the mouth of the cylinder is turned upward, 



38 Introduction to General Chemistry 

but will not escape if the mouth is downward. One liter of 
methane weighs 0.721 g., which is but little more than half of 
the weight of the same volume of air. If we bring a lighted 
candle into a cylinder of methane we find that the gas burns 
with a slightly luminous flame but that the candle flame is 
extinguished. 

55. The Quantitative Analysis of Methane. — If we examine 
the products of combustion from a methane flame we find water 
and carbon dioxide, from which we know that methane is a com- 
pound of carbon and hydrogen with or without other elements. 
We may determine the quantitative composition of methane by 
precisely the same method as that used for the quantitative 
analysis of acetylene, whereupon we find that methane contains 
75.0 per cent of carbon and 25.0 per cent of hydrogen by weight. 
Since the sum of these percentages is 100 we know that methane 
must contain only the elements carbon and hydrogen. From the 
data obtained in the analysis of methane we may also calculate 
that 1 liter of methane under standard conditions contains 
o. 180 g. of combined hydrogen. 

56. The Weight of Hydrogen in One Liter of Gaseous Hydro- 
gen Compounds. — By a study of the composition of the four 
gases, hydrogen chloride, acetylene, ammonia, and methane, as 
well as of hydrogen itself, we have found the weight of hydrogen 
in 1 liter of each. These results may now be tabulated as in 
Table II. An inspection of the results given in the table reveals 

TABLE II 

Hydrogen chloride o . 045 g. 

Hydrogen o . 090 

Acetylene o . 090 

Ammonia 0.135 

Methane o . 180 

a remarkable fact. The weight of hydrogen in 1 liter of hydrogen 
chloride is less than that in any other case. The weight per liter 
of hydrogen gas itself is double the weight of hydrogen in 1 liter of 
hydrogen chloride. Likewise the weight of hydrogen in 1 liter of 
acetylene is exactly equal to the weight of a liter of free hydrogen and 



The Law of Definite Composition 39 

also double the weight of hydrogen in 1 liter of hydrogen chloride. 
The weight of hydrogen in 1 liter of ammonia is three times that in 
1 liter of hydrogen chloride, while in the case of methane the weight 
of hydrogen per liter is four times the weight of this element in the 
same volume of hydrogen chloride. 

If we consider the weight of hydrogen in a liter of hydrogen 
chloride as unity, we find that the weights in the same volumes of the 
other gases are expressed by the numbers 2, 3, or 4. It is obvious 
that the relations we discussed would also hold equally well if 
we dealt with weights of hydrogen contained in any other fixed 
volume, as a cubic foot or a cubic meter. We could express the 
facts by saying that the weight of hydrogen contained in a fixed 
volume of any of these gases is in each case a multiple of the mini- 
mum weight, which is found in the case of hydrogen chloride gas. 
Since 1 liter of hydrogen chloride gas contains 0.045 g. of hydro- 
gen, 1 g. of combined hydrogen would be contained in 22.4 
liters 1 of hydrogen chloride. In the same volume of the other gases 
the weights of hydrogen would be 2 g., 3 g., or 4 g. 

1 In reality 1-=- 0.045 gives 22 . 2 instead of the correct value 22.4 liters. The 
discrepancy is caused by the fact that the members used are only approximate. 
This subject is discussed further in section 222. 



CHAPTER V 
SYMBOLS AND CHEMICAL FORMULAE 

57. Gaseous Carbon Compounds. — We may now inquire 
whether the remarkable relations between the weights of hydro- 
gen in equal volumes of compounds of hydrogen hold good in the 
case of compounds of other elements. We have already studied 
three gaseous compounds of carbon: carbon dioxide, acetylene, 
and methane, and have seen how the percentage composi- 
tion of each is determined. Before discussing the results so 
obtained, let us consider two new gaseous compounds of carbon : 
propane and trimethylamine. 

58. Propane: a Compound of Carbon and Hydrogen. — 
Propane is found in small amounts in the natural gas of some 
wells and also dissolved, in small quantities, in crude petroleum. 
It may also be obtained artificially by methods well known to the 
chemist, the nature of which we need not now consider. We 
observe that propane is a colorless, odorless gas which is some- 
what heavier than air, 1 liter under standard conditions weighing 
1 . 97 g. We find that propane resembles methane in its chemi- 
cal behavior, since it extinguishes a burning candle but takes 
fire itself at the same time, burning with a slightly luminous 
flame and yielding carbon dioxide and water as the only products 
of combustion. The analysis of propane may be carried out in 
precisely the same manner as our analysis of methane and acety- 
lene. We find in this way that propane contains 81 .8 per cent 
of carbon and 18.2 per cent of hydrogen. Since the sum of these 
percentages is 100, it follows that carbon and hydrogen are the 
only constituents of propane. 

59. Trimethylamine: a Compound of Carbon, Hydrogen, and 
Nitrogen. — Trimethylamine is a colorless gas about twice as 
heavy as air, 1 liter weighing 2 . 65 g. Its odor is very powerful 
and somewhat disagreeable, but if inhaled in small quantities the 
gas is not poisonous nor irritating, as is, for example, chlorine gas. 
The odor is that of decaying fish. In fact, the gas can be 

40 



Symbols and Chemical Formulae 



4i 



obtained from products separated from herring brine. We find 
that the gas is very easily soluble in water and that the solution 
turns red litmus paper blue, just as ammonia does; but the gas 
may be distinguished from ammonia by the fact that it will burn, 
whereas ammonia will not. It is easy to discover that water and 
carbon dioxide are formed when trimethylamine is burned in air 
or in oxygen. If we pass trimethylamine through a tube con- 
taining heated copper oxide we obtain, in addition to water and 
carbon dioxide, a colorless, odorless, incombustible gas which can 
easily be identified as nitrogen. These facts show that tri- 
methylamine contains the elements carbon, hydrogen, and nitrogen. 
We could determine the percentages of carbon and hydrogen by 
rinding the weights of carbon dioxide and water formed by the 
action of the gas on hot copper oxide, as in analyses previously 
made. We might also find the percentage of nitrogen by finding 
the volume of nitrogen which we could obtain from a known 
volume of the gas. The percentages of carbon, hydrogen, and 
nitrogen would be found to be 61 . o, 15.3, and 23 . 7 respectively. 
60. The Weights of Carbon in 1 Liter and in 22.4 Liters of 
Gaseous Carbon Compounds. — Let us now consider the facts 
presented in Table III. The weight of 1 liter and the percentage 

TABLE III 



Weight of 
1 Liter 



Percentage of 
Carbon 



Weight of 

Carbon in 

1 Liter 



Weight of 
Carbon in 
22.4 Liters 



Methane 

Carbon dioxide . 

Acetylene 

Propane 

Trimethylamine 



o. 72 
1.97 
1. 19 
1.97 
2.65 



75-o 
27-3 
92.3 
81.8 
61.0 



o-54 
o.54 
1.08 
1 .62 
1 .62 



12 
12 
24 
36 

30 



of carbon in each of the five gaseous compounds of carbon we 
have studied are given in the first and second columns of figures. 
The product of the weight of 1 liter of a gas by the percentage of 
carbon it contains gives the weight of combined carbon in 1 liter. 
These products are given in the third column. The weights of 
carbon in 22.4 liters, as given in the last column, are found by 
multiplying the corresponding weights in the third column by 
22.4. 



42 Introduction to General Chemistry 

We see by a glance at the last column of the table that 22.4 
liters of carbon dioxide and methane contain 12 g. of combined 
carbon, that the same volume of acetylene contains 24 g. of 
carbon, while the weight of combined carbon in 22.4 liters of 
propane and trimethylamine is 36 g., and therefore that the 
weight of carbon in 22 . 4 liters of any of these gases is either one, 
two, or three times 12 g. In the case of gaseous hydrogen com- 
pounds, we found that the weight of hydrogen was either one, 
two, three, or four times 1 g., which was the minimum weight 
of this element found in any case. We thus find that in 22.4 
liters of various pure gases the minimum weight of hydrogen is 1 g. 
and the minimum weight of carbon 12 g., and, further, that if a 
greater weight of either of these elements is contained in this volume 
of any pure gas, the weight is a multiple of the minimum weight 
by a small whole number. 

Let us now consider the weights of carbon and hydrogen 
contained in 22.4 liters of the three gaseous compounds which 
contain only carbon and hydrogen, namely, methane, acetylene, 
and propane. In 22.4 liters of methane we find 12 g. of carbon 
combined with 4 g. of hydrogen. In the same volume of acety- 
lene, 24 g. of carbon combined with 2 g. of hydrogen, and in the 
case of propane 36 g. of carbon combined with 8 g. of hydrogen. 
Without considering at present the theoretical significance of 
the remarkable facts which these figures show, we may consider 
a practical application of the facts which will enable us to express 
the composition of these gases in a simple fashion. 

The student must realize that since we have three compounds 
all consisting of carbon and hydrogen and having different prop- 
erties, the difference in percentage composition must be an 
important factor in determining the properties of the substance. 
He will also understand that a knowledge of the percentage com- 
position is a matter of prime importance for the chemist, and that 
any scheme by means of which a knowledge of the composition 
by weight could be easily memorized would be important. 

61. Symbols. — Suppose we represent 1 g. of hydrogen by a 
sign or symbol and choose the letter H for this purpose. We 
could, then, represent by H taken four times the weight of hydro- 



Symbols and Chemical Formulae 43 

gen contained in 22.4 liters of methane; by H taken twice, or 
2H, the amount of hydrogen in 22.4 liters of acetylene; and 
similarly by 8H, the amount of hydrogen in 22 . 4 liters of propane. 
Suppose that, on the other hand, we represent 12 g. of carbon by 
the sign or symbol C, then C, 2C, and 3C will represent the 
weights of carbon in 22.4 liters of methane, acetylene, and 
propane respectively. The weights of carbon and hydrogen in 
22.4 liters of methane may then be represented by writing iC 
together with 4H. As a matter of convenience the multiples 1 
for the C and 4 for the H, are written as subscripts; so that 
instead of iC and 4H we write dH 4 . In practice no subscript 
is used when the multiple is 1. The composition of methane is 
represented simply by CH 4 . 

62. Chemical Formulae. — In a similar way we may represent 
the weights of carbon and hydrogen in 2 2 . 4 liters of acetylene by 
C 2 H 2 while the composition of the same volume of propane may 
be represented by C 3 H 8 . We call H the symbol for hydrogen, and 
for the present we may consider that H or iH represents 1 g. 
of hydrogen and similarly that C, the symbol for carbon, repre- 
sents 12 g. of that element. We call the expressions CH 4 , C 2 H 2 , 
and C 3 H 8 the formulae of methane, acetylene, and propane 
respectively. We shall now proceed to show how this system 
may be extended to all gaseous compounds of any element what- 
ever. 

Chemists are familiar with a large number of gases in addi- 
tion to those which we have already studied. Some of these are 
of much practical importance while others are chiefly of interest 
to the chemist for scientific reasons. In every case it is a simple 
matter to determine the weight of 1 liter of the gas under stand- 
ard conditions, the method of making the determination being 
essentially the same in all cases. Furthermore, by methods 
which are well known to chemists we may determine what ele- 
ments compose any gas, and by means of a quantitative analysis 
we may determine the percentage of each element in the gas. 
If we calculate in the case of each gas the weight of each element 
contained in 22.4 liters of the gas, we obtain results like those 
shown in Table IV. 



44 



Introduction to General Chemistry 



63. The Minimum Weights of Oxygen, Nitrogen, and 
Chlorine. — An inspection of the results given in Table IV shows 
that the same regularity in the weights of hydrogen and carbon 
holds in all cases, as we have observed it to hold in the few cases 
discussed in the preceding paragraphs. We notice also that the 
minimum weight of oxygen in 22.4 liters of any of its gaseous 

TABLE IV 

Weights of Constituents in 22.4 Liters of Gases 



Substance 



Oxygen _ 

Carbon monoxide . . 
Carbon dioxide .... 

Nitrous oxide 

Nitric oxide 

Nitrosyl chloride . . . 
Hypochlorous oxide. 
Chlorine dioxide . . . 

Phosgene 

Methyl ether 

Hydrogen 

Hydrogen chloride. . 

Prussic acid 

Ammonia 

Methane 

Acetylene 

Ethylene 

Ethane 

Propylene 

Propane 

Methyl chloride. . . . 

Ethyl chloride 

Methylamine 

Nitrogen 

Cyanogen 

Cyanogen chloride. . 

Chlorine 

Trimethylamine . . . . 



Oxygen 



2X16 
1X16 
2X16 
1X16 
1X16 
1X16 
1X16 
2X16 
1X16 
1X16 



Hydrogen 



6X1 

2X1 

1X1 
1X1 

3X1 
4X1 
2X1 
4X1 
6X1 
6X1 
8X1 
3X1 
5X1 
5X1 



9X1 



Carbon 



IXI2 
IXI2 



IXI2 
2Xl2 



IXI2 



1X12 
2X12 
2X12 
2Xl2 
3X12 
3X12 
IXI2 
2X12 
IXI2 



2X12 
IXI2 



3Xl2 



Nitrogen Chlorine 



2X14 
1X14 
1X14 



1X14 
1X14 



1X14 
2X14 
2X14 
1X14 



1X14 



1X35-5 
2X35-5 
1X35-5 
2X35.5 



1X35-5 



1X35-5 
1X35-5 



1X35-5 
2X35-5 



Formula 



2 

CO 

C0 2 

N 2 

NO 

NOC1 

C1 2 

C10 2 

COCl 2 

C 2 H 6 

H 2 

HC1 

HNC 

NH 3 

CH 4 

C 2 H 2 

C 2 H 4 

C 2 H 6 

C3H6 

C 3 H 8 

CH3CI 

C 2 H S C1 

CH 5 N 

N 2 

C 2 N 2 

C1NC 

Cl 2 

C 3 H 9 N 



compounds is 16 g., and that this weight is found in many 
cases, while in others the weight is twice 16. In the case of 
the compounds of nitrogen we note that the minimum weight 
is 14 g. and that in other cases the weight is double this mini- 
mum weight. In the case of chlorine compounds the minimum 
weight of chlorine is 35.5 g., while those compounds with 
a larger proportion of chlorine contain double the minimum 
weight. 



Symbols and Chemical Formulae 45 

64. The Law of Minimum and Multiple Weights. — Entirely 
analogous regularities will be found if we consider the data 
obtained from a study of the gaseous compounds of any other 
elements. For each element we find a minimum weight in the 
volume of 22 .4 liters of any of its gaseous compounds under standard 
conditions and also find that the weight if greater than the minimum 
would be 2, 3, or 4, or some small multiple of this minimum. This 
last statement may be called the Law of Minimum and Multiple 
Weights. 

65. The Chemical Unit Volume: 22.4 Liters. — The volume 
22.4 liters thus becomes a kind of unit volume for the chemist, 
this particular volume having been chosen because it contains 
1 g. of hydrogen in the case of those hydrogen compounds 
which contain the minimum weight of this element. In this 
volume no other element has a minimum weight as small as that 
of hydrogen. 

66. Symbols Represent Minimum Weights. — In the same 
manner as that suggested in a preceding paragraph for hydrogen 
and carbon, we may represent the minimum weight of each of 
the other elements by a symbol. Table V shows the minimum 
weights of the five elements we have been considering, together 
with the corresponding symbols. 



Minimum Weights 


TABLE 

IN 22.4 


V 

Liters, and 


Symbols 


Hydrogen 

Carbon 


I.Og. 

12.0 
14.0 
16.0 

35-5 


H 
C 


Nitrogen 

Oxygen 


N 



Chlorine 


CI 







67. Making Formulae. — We see from Table IV that 22.4 
liters of carbon dioxide contain 12 g. of carbon combined with 
2X16 g. of oxygen. We may, therefore, represent the composi- 
tion of the quantity of carbon dioxide in 22.4 liters by the 
formula C0 2 . In an analogous fashion we may obtain as the 
formula representing the composition of 22.4 liters of ammonia, 
NH 3 , and as the formula for hydrogen chloride, HC1. By making 



46 Introduction to General Chemistry 

use of this system the student will now have no difficulty in 
writing down at once the formula of each of the gases from the 
data contained in Table IV. He will also readily see that it is a 
much less difficult task to learn the formulae of such gases than 
to learn their percentage composition; that is to say, it is an easier 
tax upon the mind to remember the formula HC1 than to remem- 
ber that hydrogen chloride contains 2 . 76 per cent of hydrogen 
and 97. 24 per cent of chlorine. 

68. The Practical Use of Formulae. — A review of the methods 
employed in arriving at the results represented by the formula 
of any substance shows that in each case we have made use of the 
knowledge, first, of the weight of 1 liter of the gas, and, secondly, 
of the percentage of each of its elementary constituents. Con- 
versely, if we know the weights which the symbols of the elements 
represent, and know the formula of a gas, we may by working 
backward find its percentage composition. For example, sup- 
pose that we remember that the formula of methane is CH 4 
and know that H stands for 1 g. of hydrogen and C for 12 g. 
of carbon. Then 22.4 liters of methane contain 4 g. hydrogen 
combined with 12 g. of carbon. The proportion of hydrogen is, 
therefore, 4/16, or 25 percent, and of carbon 12/16, or 75 per cent, 
the weight of 22 . 4 liters being 16 g., and 1 liter weighs 16/22.4 = 
0.72 g. In calculating in this way the density and percentage 
composition of methane we are merely reproducing the results 
which originally were obtained by experiment. In order to find 
the formula of any gas, we must know its density and the per- 
centage of each elementary constituent. We find by actual 
experience that we can represent by a formula, usually of a very 
simple character, the composition of 22.4 liters of any gaseous 
substance. 

69. Formulae of Liquids and Solids. — The system which we 
have just considered is capable of extension to liquid and solid 
substances, in which case, however, the formula may have a 
slightly less definite meaning. We may illustrate this by con- 
sidering the cases of water and mercury oxide. We have found 
that water is composed of 1 1 . 2 per cent of hydrogen and 88 . 8 
per cent of oxygen, from which we observe that the weight of 



Symbols and Chemical Formulae 47 

the oxygen is 8 times the weight of the hydrogen with which it is 
united. This ratio of hydrogen and oxygen might be represented 
by H 2 0, since this formula would mean that 2 g. of hydrogen are 
united with 16 g. of oxygen, which weights of hydrogen and 
oxygen are in the ratio of 1 to 8, but the formulae H 4 2 and 
H6O3 would also represent equally well the proportion of hydro- 
gen and oxygen actually found in water. 

70. The Formula of Water. — We may be led to choose a con- 
sistent formula for water by the consideration of the density of 
water vapor or steam; but in this case the density determination 
must be made at a temperature above the boiling-point of water, 
if we work at atmospheric pressure. Since the effect of changes 
of pressure and temperature upon the volume of a given quantity 
of steam are the same as upon an equal volume of any gas which 
would not liquify if cooled to o° at 76 cm. pressure, we might 
calculate by the laws of Boyle and Charles what the volume of 
the known weight of steam measured at a high temperature and 
known pressure would be if the steam were under standard con- 
ditions, that is, at o° and 76 cm. pressure.- It has been found in 
this way that 1 liter of water vapor if it did not condense to a 
liquid would weigh o . 806 g. under standard conditions, which 
corresponds to a weight of 18 g. for 22.4 liters. Now, 11 . 2 per 
cent of 18 g. is 2 g. and 88.8 per cent of 18 g. is 16 g. From 
these results we conclude that if water vapor could exist under 
standard conditions as a gas that 22.4 liters would contain 
2 g. of combined hydrogen and 16 g. of combined oxygen, which 
amounts would be exactly represented by the formula H 2 0. 

71. Formulae of Volatile Liquids and Solids. — In a perfectly 
analogous fashion we could find the formula for any other volatile 
substance, the density of whose vapor we could measure experi- 
mentally. Such a procedure would enable us to represent by a 
formula the composition of a great number of volatile chemical 
substances which are not gaseous, but are liquid or solid under 
ordinary conditions of temperature and pressure. 

72. Formulae of Involatile Substances. — There are, however, 
many chemical substances which are not volatile or which can- 
not be volatilized at temperatures at which we could make 



48 Introduction to General Chemistry 

experimental determinations of their vapor densities. There 
are other solids and liquids which would be decomposed if 
strongly heated. For such substances we could not find chemi- 
cal formulae in the same way as for gases or volatile substances. 
However, we can and do represent by formulae the composition 
of such involatile substances. 

73. The Formula of Red Oxide of Mercury. — The method of 
obtaining the formula of such a substance may be illustrated by 
the case of the red oxide of mercury, which, it will be remem- 
bered, is readily decomposed when heated into mercury and 
oxygen. We find by analysis that this compound contains 
92.6 per cent of mercury and 7.4 per cent of oxygen. By the 
experimental study of volatile mercury compounds, as well as of 
mercury itself, we find that the minimum weight of mercury in 
22.4 liters is 200 g., and therefore represent this weight of 
mercury by the symbol Hg. It now remains to discover what 
multiples of 200 for the mercury and of 16 for the oxygen are 
in the same ratio as the percentages of mercury and oxygen in 
mercury oxide. We find very easily that 200 is to 16 as 92.6 is 
to 7.4, and from this we write the formula HgO. 

We could of course represent the same proportions of mercury 
and oxygen by the formula Hg 2 2 . But we are not able to 
decide which of these to choose as in the case of a volatile sub- 
stance where the formula represents the quantity of material in 
22.4 liters of the gas or vapor under standard conditions. In 
such a case we choose the simpler formula, in this case HgO, but 
we must bear in mind that the formula does not mean quite as 
much in such a case as in that of a gas or volatile substance, 
where it always represents in addition to the true proportion 
of the constituent elements the actual weights of each in 22.4 
liters of the gas under standard conditions. 

74. Symbol Weights and Formula Weights. — The letter or 
pair of letters which represents the minimum weight of an 
element in 22.4 liters of any of its gaseous compounds is called 
the symbol of that element and the weight which this symbol 
represents may then be called the symbol weight. Each of the 
eighty-five or more known elements has been assigned a definite 



Symbols and Chemical Formulae 49 

symbol which represents a definite symbol weight. We have 
seen (62) how the quantities of each element in 22.4 liters of a 
compound gas may be represented by a formula made up of 
symbols, each symbol being multiplied by a factor which shows 
how many times the minimum weight of the element is present 
in 22 . 4 liters of the gaseous compound. The sum of the weights 
represented by the various symbols each multiplied by its factor 
is naturally the weight of 22.4 liters of the gas, represented by 
the formula. This weight is often spoken of as the formula 
weight. In the case of an involatile solid substance the formula 
weight is the -weight represented by the formula but indicates 
only theoretically the weight which we should expect 22.4 liters 
of the substance to have if it were a gas under standard con- 
ditions. 

75. The Formulae of Some Elementary Gases. — It is impor- 
tant to note that 22.4 liters of the gases hydrogen, oxygen, 
nitrogen, and chlorine weigh 2, 32, 28, and 71 g. respectively 
(63, Table IV). These weights are for each element just double 
the minimum weights which we find in numerous compounds 
of the elements and therefore in each case just double the weight 
represented by the symbol. We must therefore write, as the 
formulae of these gases, H 2 , 2 , N 2 , and Cl 2 , respectively. The 
formula of an elementary gas in the free state will then represent 
the quantity of that gas in 22.4 liters. We must here point 
out that not every element in the form of gas or vapor is to be 
represented by a formula composed of its symbol taken twice. 
For example, the vapors of mercury and sodium have the single 
symbol formulae Hg and Na, respectively; on the other hand, 
the formulae of the vapors of the elements phosphorus and sulfur 
are P 4 and S 8 . 



CHAPTER VI 
CHEMICAL EQUATIONS 

76. Equations. — In this chapter we shall see how it is possible 
to represent in a very simple way the quantities of substances 
entering into and formed in a chemical reaction. Let us con- 
sider the case of hydrogen and chlorine which has already been 
studied experimentally. We have learned that hydrogen and 
chlorine unite to form hydrogen chloride (44). Furthermore 
we find by experiment that one volume of hydrogen and one 
volume of chlorine give two volumes of hydrogen chloride; so 
that if 22.4 liters of hydrogen united with 22.4 liters of chlorine 
we should obtain 44.8 liters of hydrogen chloride. Now we 
may represent 22.4 liters of hydrogen by the formula H 2 and 
22.4 liters of chlorine by Cl 2 , while for twice 22.4 liters of hydro- 
gen chloride we put the coefficient, 2, in front of the formula and 
write 2HCI. We may then express the facts by stating that H 2 

plus Cl 2 gives 2HCI or 

H 2 +C1 2 ->2HC1 

which may also be written 

H 2 +C1 2 =2HC1. 

We call either of these expressions the equation for the reaction 
between hydrogen and chlorine. 

77. What an Equation Means. — The equation 

H 2 -f-Cl 2 -^ 2 HCl 

expresses the fact that the quantity of hydrogen represented by 
the formula H 2 or 2 g. unites with the quantity of chlorine 
represented by Cl 2 or 71 g. to give the quantity of hydrogen 
chloride represented by 2HCI or 73 g. It also expresses the 
fact that 22.4 liters of hydrogen unite with 22.4 liters of chlorine 
to give 2X22.4 liters of hydrogen chloride, or in general that one 
volume of hydrogen and one volume of chlorine unite to give 
two volumes of hydrogen chloride, the volumes being those of 



Chemical Equations 51 

the gases measured in all cases under standard conditions. In 
reactions involving gases the volume of each gas taken or 
formed is always shown by the coefficient in front of its formula 
in the equation for the reaction. 

78. The Equation for the Burning of Carbon. — Some free 
elements like carbon are not sufficiently volatile to enable us to 
find the formula of the element from measurements of the vapor 
density of the free element, and in such a case we use the symbol 
of the element in writing equations involving its reactions. 
When carbon is burned we find that 12 g. of carbon require 32 g. 
of oxygen occupying a volume of 22.4 liters, and producing 44 g. 
of carbon dioxide occupying also a volume of 22.4 liters. These 
facts may therefore be represented by the equation 

c+o 2 ->co 2 . 

Here the equation expresses directly the weights of carbon and 
oxygen which unite as well as the weight of carbon dioxide 
formed. At the same time it also shows that 22.4 liters of 
oxygen when completely combined with sufficient carbon gives 
22.4 liters of carbon dioxide, but since the carbon is not in the 
gaseous state the equation does not indicate anything regarding 
the volume of the solid carbon which unites with the volume of 
oxygen represented by the formula 2 . 

79. Solving Problems. — If we remember that the equation 
for the burning of carbon in oxygen is 

C+0 2 ->C0 2 

we may make use of the facts represented by the equation in the 
solution of problems such as the following: How many liters 
of oxygen are required for the burning of 5 g. of carbon? To 
solve this problem we first write down the equation which repre- 
sents the reaction. This shows that the quantity of carbon 
represented by the symbol C, namely, 12 g., requires for its com- 
bustion the volume of oxygen represented by the formula O a , 
namely, 22.4 liters. Therefore 5 g. of carbon would require 
the volume determined by the proportion 

12:5: :22 .4::v 






52 Introduction to General Chemistry 

where x is the number of liters of oxygen necessary for the com- 
bustion of 5 g. of carbon. In an analogous manner we may 
calculate what volume of carbon dioxide is produced by the 
burning of a known weight of carbon. 

We may also calculate what weight of oxygen is required or 
carbon dioxide produced in the burning of 5 g. of carbon. If 
12 g. of carbon require 32 g. of oxygen, as our equation indicates, 
then we have only to solve the following proportions in order to 
find the weight of oxygen required for 5 g. of carbon: 

i2:5::32:y 

where y is the required answer. 

80. The Burning of Magnesium. — Suppose we desire to find 
by experiment the formula of the product formed by burning 
magnesium in oxygen. It will ,be recalled that the metal 
magnesium in the form of wire or powder burns with great ease 
in oxygen, forming a white solid substance which we have called 
magnesium oxide (28). We find by experiment that 10 g. of 
magnesium when burned yields 16.6 g. of magnesium oxide. 
Let us suppose that we have discovered by careful experiment 
that magnesium oxide contains only the elements magnesium 
and oxygen. The difference between the weight of the mag- 
nesium oxide formed and the magnesium taken must represent 
the weight of oxygen which has combined with the 10 g. of 
magnesium. This we find to be 6 . 6 g. 

Suppose we know that the symbol weight of magnesium is 
24.3 g. or 

Mg = 24.3g. 

It is now required to calculate the relative numbers of symbol 
weights of magnesium and oxygen that unite to form magnesium 
oxide. We know that 10 g. of magnesium unite with 6.6 g. of 
oxygen. We may then make the proportion 

10:6.6: 124.3:2 

from which we find that 2=16. Therefore 16 g. of oxygen 
represented by 0, combine with the weight of magnesium repre- 



Chemical Equations 53 

sented by the symbol Mg, and consequently we may represent 
the composition of magnesium oxide by the formula MgO and 
write the equation for the burning of magnesium thus: 

Mg+0->MgO 
or better 

2Mg+0 2 ->2MgO, 

the latter equation having the advantage in that it shows the 
volume of oxygen, 22.4 liters, as well as its weight required for 
the burning of the weight of magnesium represented by 2Mg. 
But since both magnesium and magnesium oxide are solid 
involatile substances the equation does not show the volumes of 
these solids entering into the reaction, as it would in the case of 
gaseous substances. 

81. The Burning of Iron. — It will be recalled (17) that iron 
burns in oxygen, giving iron oxide, the formula for which we 
may now calculate. In an experiment in which 12.6 g. of iron 
was burned the weight of iron oxide produced was 17.4 g., from 
which we find, by subtracting the weight of the iron burned, the 
weight of the oxygen to be 4.8 g. These weights of iron and 
oxygen must be in the same ratio that some number of times 
56, the symbol weight of iron, is to some number of times 16 
where these numbers are small integers. Dividing 12.6 by 56 
we get 0.225. Dividing 4.8 by 16 we get 0.300. Since these 
numbers 0.225 an d 0.300 are not equal, the formula cannot 
be FeO. It will, however, readily be found that 0.225 is to 
0.300 as 3 is to 4, and therefore that 12.6:4.8: 13X56:4X16, 
which shows that the formula of the oxide of iron formed by 
burning iron in oxygen is Fe 3 4 . We may then write the equa- 
tion for the burning of iron as follows : 

3Fe+20 2 ->Fe 3 4 . 

82. The Action of Hydrogen on Copper Oxide.— It will be 
remembered that we found earlier that heated copper oxide and 
hydrogen give metallic copper and water (33). In a quantitative 
experiment it was found that 2.387 g. of copper oxide yielded 
1.907 g. of copper and o.54g. of water. From the weights of 



54 Introduction to General Chemistry 

copper and Gopper oxide, together with a knowledge of the fact 
that copper oxide is composed of copper and oxygen only, we may 
discover very readily that the formula of copper oxide is CuO, 
knowing the symbol weight of copper to be 63 . 6. Furthermore, 
since water contains only hydrogen and oxygen and 0.54 g. of 
water has been formed from 2.387—1.907 or 0.48 g. of oxygen, 
the weight of hydrogen present in the o . 54 g. of water must have 
been o . 06 g. Making a calculation analogous to that made in 
finding the formula for iron oxide, we find that 0:06:0.48: : 2X 
1:1X16 and that therefore the composition of water is repre- 
sented by the formula H 2 0. We may now write, as the equation 
for the reaction between copper oxide and hydrogen, 

CuO+H 2 ^Cu+H 2 0. 

83. The Action of Acetylene on Copper Oxide. — From what 
has preceded the student will understand that in order to be able 
to write the equation for any reaction we must know all of the 
substances entering into the reaction and all of the products. In 
addition we must know the formula of each substance. We may 
illustrate the method then employed by means of reaction 
between acetylene and copper oxide which we have already 
studied. 

When acetylene is passed over heated copper oxide we obtain 
carbon dioxide and water, while metallic copper is left behind, 
these three substances being the sole products of the reaction (50) . 
The formula of acetylene is C 2 H 2 (62) . The quantity of carbon 
represented by C 2 would give the quantity of carbon dioxide 
represented by 2C0 2 ; and the quantity of hydrogen represented 
by H 2 would give the quantity of water represented by H 2 0, so 
that the quantities of carbon and hydrogen represented by one 
formula weight of acetylene C 2 H 2 would yield the quantities of 
carbon dioxide and water represented by 2C0 2 +H 2 0. The 
quantity of oxygen contained in the quantities of carbon dioxide 
and water represented by 2C0 2 +H 2 is represented by 5O, 
which quantity is contained in the amount of copper oxide 
represented by 5 CuO. It will thus appear that the quantity 
of acetylene represented by C 2 H 2 will require the quantity of 



Chemical Equations 55 

copper oxide represented by 5C11O, and there will be produced 
the quantities of the three products represented by 

2 C0 2 +H 2 0+5Cu. 

The equation is therefore 

C 2 H 2 +5CuO-»2C0 2 +H 2 0+5Cu. 

84. The Action of Ammonia on Copper Oxide. — In an analo- 
gous manner we may obtain as the equation for the reaction 
which occurs when ammonia gas is passed over heated copper 
oxide, in which case water, nitrogen, and metallic copper are 

formed, 

2NH 3 + 3 CuO->3H 2 0+3Cu+N 2 . (52) 

85. The Meaning of an Equation. — Since chemists make 
extensive use of equations, it is of fundamental importance that 
the student should understand exactly how equations are ob- 
tained and what they mean. In every case before the equation 
for the reaction can be written the reaction must have been 
thoroughly investigated by experiment in the manner illustrated 
in the preceding examples. The equation then shows at a glance 
what substances enter into and are formed as a result of the 
reaction. It also shows the composition of each of the substances 
concerned and the proportions in which they take part in the 
reaction, it being assumed in all cases that we know the weight 
for which the symbol of each element stands. 

86. An Equation Balances. — It is one of the most funda- 
mental facts in chemistry that in chemical change no material is 
destroyed but that the elements merely change their forms of 
combination with one another. This important fact, which we 
know as the Law of the Indestructibility of Matter, is also repre- 
sented in every chemical equation. For we notice that in each 
equation we have on each side the same number of symbol 
weights of each element. Thus in the equation 

C 2 H 2 + 5 CuO->2C0 2 +H 2 0+5Cu 

we see that there are on each side two symbol weights of carbon, 
two symbol weights of hydrogen, five symbol weights of copper. 



56 Introduction to General Chemistry 

and five symbol weights of oxygen. This fact is usually expressed 
by saying that the equation balances. 

All of the reactions which we have studied up to this time 
have been thoroughly investigated by chemists and for each the 
reaction equation has been discovered. We may now give, in 
Table VI, a list of such equations for purposes of reference. It 
is not to be expected, however, that the student should make 
great effort to memorize all of these equations, although such a 
task would not be very difficult, for, as a little inspection will 
show, there are certain regularities observable which make this 
a less difficult task than might at first sight seem to be the case. 

TABLE VI 
Equations of Other Reactions Studied 
2Hg+0 2 -> 2 HgO 
2 H 2 +0 2 ->2H 2 
2Na+2H 2 0->2NaOH+H 2 
XaOH+HCl->NaCl+H 2 
2Na+Cl 2 ->2NaCl 
Mg+H 2 0->MgO+H 2 
3 Fe+4H 2 0->Fe 3 4 +4H 2 
CH 4 +20 2 ->C0 2 +2H 2 
C 3 H 8 +ioCuO->3C0 2 +4H 2 0+ioCu 

87. Problems 

1. What weight of mercury can be obtained by the decom- 
position of 10 g. of mercuric oxide? 

2. Wnat volume of oxygen at o° and 76 cm. can be made 
from 8 g. of mercuric oxide ? 

3. What weight of sodium must be acted on by water to 
yield 500 c.c. of hydrogen at o° and 76 cm.? 

4. What weight of common salt can be made from 10 g. of 
metallic sodium? 

5. What volume of hydrogen at 20 and 72 cm. would be 
formed by the action of sufficient steam on 6 g. of magnesium? 

6. What weight of copper oxide would be required for the 
oxidation of 200 c.c. of propane measured at 25 and 74 cm.? 
(See last equation of Table VI above.) 

What weight of water would be formed? 



CHAPTER VII 
ACIDS, BASES, AND SALTS— I 

88. Caustic Soda or Sodium Hydroxide. — Let us now con- 
sider the chemical changes which occurred in the formation of 
common salt from metallic sodium, which we have already 
studied experimentally. It will be recalled that sodium reacted 
violently with water, giving hydrogen and sodium hydroxide, the 
reaction being represented by the equation 

2 Na+ 2 H 2 -> 2 NaOH+H 2 . (40) 

If we repeat the experiment and evaporate the water we find 
that sodium hydroxide (also known as caustic soda) is left as a 
white solid which is readily soluble in water. This solution feels 
" soapy" to the fingers and if greatly diluted with water is found 
to have an unpleasant "soapy" taste. (It must not be tasted 
unless greatly diluted with water, since the concentrated solu- 
tion acts powerfully on the mucous membrane.) A piece of red 
litmus paper is turned blue if dipped in the solution. We know 
many other substances which have properties similar to those 
of sodium hydroxide. Such substances are called bases; they 
also have other characteristic properties, the most important of 
which we may now consider. 

89. Bases Neutralize Acids. — We have learned (41) that 
caustic soda and hydrochloric acid (which is a solution of hydro- 
gen chloride in water) react to give common salt. The equation 
for this reaction is 

NaOH+HCl-> H 2 0+NaCl. 

If we add more than sufficient of the acid and then evaporate 
the solution to dryness, the excess of hydrogen chloride will pass 
off with the water and nothing but pure salt, the chemical name 
of which is sodium chloride, will remain. If we test hydro- 
chloric acid with blue litmus we find that the latter is turned red. 
even by a very dilute solution. But we find that a solution of 

57 



58 Introduction to General Chemistry 

pure common salt in water has no effect on either blue or red litmus: 
it is neutral. 

90. Properties of Acids. — If we again add, drop by drop, a 
solution of hydrogen chloride to one of sodium hydroxide to 
which a few drops of a solution of litmus have been added, we 
find that the change of color from blue to red is produced suddenly 
and not gradually, a single drop being sufficient to cause the 
change. If we stop adding hydrogen chloride at this point we 
find that the solution consists only of pure salt and water (with 
but a minute amount of litmus). It no longer has the taste of 
the sodium hydroxide, but only that of salty water. A diluted 
solution of hydrogen chloride has a rather agreeable sour taste, 
reminding one of vinegar or lemon juice. Our experiment has 
shown that both the taste and the behavior toward litmus of 
sodium hydroxide and hydrogen chloride have been changed in 
their interaction. We say that they have neutralized each other. 
We know very many substances which will neutralize sodium 
hydroxide; all of these have a sour taste and color litmus red. 
We call such substances acids, the common name of hydrogen 
chloride solution being hydrochloric acid. 

91. Another Base; Ammonium Hydroxide. — As we have 
already seen (51), ammonia gas dissolves readily in water, giving 
a solution which turns litmus blue, and we are not surprised to 
find that it neutralizes hydrochloric acid. If we evaporate the 
neutralized solution we obtain a white crystalline substance, the 
composition of which is represented by the formula NH 4 C1. 
Since ammonia gas has the formula NH 3 and hydrogen chloride 
the formula HC1, we might be inclined to write the equation 

NH 3 +HC1->NH 4 C1, 

and, in fact, just this reaction takes place if we bring the two 
gases together, a dense white cloud of the solid product being 
formed. However, if a very concentrated solution of ammonia 
in water is cooled to a very low temperature, we may obtain 
crystals of a substance having a composition represented by the 
formula NH 4 OH and called ammonium hydroxide. This sub- 
stance is formed thus: 

NH 3 +H 2 0->NH 4 OH. 



Acids, Bases, and Salts — / 59 

We might think to obtain it by the evaporation of the water 
solution of ammonia; but instead we get only ammonia gas and 
water vapor. In fact, the crystals of ammonium hydroxide 
obtained at a low temperature undergo a similar change if they 
are not kept very cold. We say that ammonium hydroxide 
dissociates readily into ammonia and water. Chemists think that 
in a water solution of ammonia part of the latter is combined 
with water to form ammonium hydroxide. It is this substance 
which is thought to act directly on red litmus, changing it to 
blue, and to act on hydrochloric acid as follows: 

NH 4 OH+HCl-> NH 4 C1+H 2 0. 

We therefore call ammonium hydroxide a base. 

92. Ammonium Chloride, Salts. — The substance NH 4 C1 is 
called ammonium chloride. In appearance, taste, and other 
properties to be studied later, sodium chloride and ammonium 
chloride closely resemble one another. They are examples of an 
important class of chemical substances called salts. 

A review of the two neutralizations just discussed will show 
that they have much in common: in each case a base reacts with 
an acid to form a salt and water. Somewhat later, other important 
facts regarding neutralization will be discovered. Before dis- 
cussing such matters we will first become acquainted with a few 
other important acids, bases, and salts. 

93. Sulfuric Acid. — One of the most important, if not the 
most important, of all acids is a substance which is known as oil 
of vitriol or sulfuric acid. It is manufactured in immense 
quantities and is very cheap, the commercial grade selling for 
less than one cent a pound. We shall not now consider the 
method of its manufacture further than to state that it is made 
from sulfur. Its composition is represented by the formula 
H 2 S0 4 . It is a colorless liquid of "oily" consistency, but is 
not really an oil, as it will mix with water in all proportions. 
It must be handled with caution, since it can cause bad burns 
if it is spilled on the skin. (In case of accident, wash off the acid in 
much running water, immediately.) When sulfuric acid is mixed 
with water, the mixture gets boiling hot, for which reason the acid 



60 Introduction to General Chemistry 

should be added very slowly, with stirring, to the water, if a 
dilute solution is to be made. 

94. Neutralization of Sulfuric Acid, Sodium Sulfate. — We 
find that the dilute solution has a sour taste and that it turns 
litmus red. We may next try whether it will neutralize a solu- 
tion of sodium hydroxide, for which purpose we may add to a 
dilute solution of sulfuric acid a few drops of litmus solution and 
then run in sodium hydroxide solution drop by drop until neutral- 
ity is reached. If the neutral solution is now boiled until a solid 
begins to appear and then is left to evaporate at room tempera- 
ture, large, transparent, glassy-looking crystals will be formed. 
These crystals dissolve readily in water to form a neutral solu- 
tion, which does not have a sour taste. 

If we allow the dry crystals to remain in the open air we find 
that they lose weight rapidly and turn white upon the surface, 
forming a fine white powder. Finally nothing is left of the large, 
clear, glassy crystals; only the powder remains, the weight of 
which is much less than that of the original material. What is 
the cause of this curious change? Let us put one of the large 
clear crystals into a dry test tube and heat gently the lower end 
of the tube containing the crystal, while the tube is held nearly 
horizontally. We soon see that water has collected in large 
amount in the cold end of the tube, while only a white powder 
is left behind. It is now easy to understand what occurred when 
the large crystal was exposed in the open air. It dissociated into 
the white powder and water which disappeared as vapor. The 
analysis of the thoroughly dried powder would show that it 
contains only sodium, sulfur, and oxygen, and in the proportions 
represented by Na 2 S0 4 , and since the clear crystals yielded only 
Na 2 S0 4 and water, their composition must be represented by 
Na 2 S0 4 . a;H 2 0, where x is a whole number which must be found 
by means of a quantitative analysis. We call the original sub- 
stance the hydrate of sodium sulfate, a hydrate of a salt being 
a compound of the salt with water. 

We may now make the equation for the formation of this 
salt from sulfuric acid. We took H 2 S0 4 and NaOH and got 
Xa 2 S0 4 , from which we see that if two formula weights of water 



Acids, Bases, and Salts — / 61 

were formed from one formula weight of H 2 S0 4 and two of NaOH, 
the whole of the material taken would be accounted for thus : 
H 2 S0 4 + 2 NaOH -> Na 2 S0 4 + 2H 2 0. 

This conclusion is rendered probable by the fact that in the other 
neutralizations we have studied water was always one of the 
products; it may be confirmed by mixing with dry sodium 
hydroxide pure sulfuric acid, whereupon water and Na 2 S0 4 will 
result. The salt Na 2 S0 4 is called sodium sulfate. Crystals of 
anhydrous sodium sulfate are different in form from those of the 
hydrate. 

95. Quantitative Analysis of a Hydrate. — Let us now consider 
the quantitative composition of the large, glassy crystals which 
yielded Na 2 S0 4 and water. If we weigh a crystal contained in a 
porcelain dish and allow it to stand a day or two at room tempera- 
ture we find that only the white powder remains. If we now heat 
the dish and contents over a flame in order thoroughly to dry 
the powder, and let it cool and weigh it again, it is obvious that 
the loss of weight will represent the weight of water originally 
combined with the weight of dry Na 2 S0 4 left in the dish. 

96. Sodium Sulfate Decahydrate; Na 2 S0 4 -ioH 2 0. — Now 
suppose that 5 . 796 g. of the hydrate of sodium sulfate yielded 
2.556 g. of dried sodium sulfate, Na 2 S0 4 , what is the formula of 
the hydrate? In other words, what is the numerical value of 
x in the formula Na 2 S0 4 -xH 2 0? The weight of water driven off 
was 5.796— 2.556 = 3.240 g. We may therefore write the 
proportion, 2.556 is to 3.240 as the formula weight of sodium 
sulfate is to the x times the formula weight of water. Now, the 
formula weight of sodium sulfate is 2X23+32+4X16 = 142 
and that of water is 2X1 + 16 = 18. Therefore 2.556:3.240:: 
142: i8x, from which we find that x = 10, and are thus led to the 
conclusion that the hydrate of sodium sulfate has the formula 
Na 2 S0 4 '101120. If the reaction between sulfuric acid and 
sodium hydroxide is represented by the equation H 2 S0 4 + 
2NaOH->Na 2 S0 4 +2H 2 0, then the hydrate Na 2 S0 4 -ioH 2 
must have resulted from the union of the sodium sulfate with 
part of the water which formed the solution, thus : 

Na 2 S0 4 + ioH 2 -> Na 2 S0 4 . 10ILO. 



62 Introduction to General Chemistry 

This substance is called sodium sulfate decahydrate (deca mean- 
ing ten). 

97. Hydrates. — Sodium sulfate forms other compounds 
with water, namely Na 2 S0 4 «7H 2 and Na 2 S0 4 «H 2 0; but the 
decahydrate is the common one. Many other salts form 
hydrates and some form a series of hydrates, as this salt does. 
But it must not be supposed that all salts form hydrates. For 
example, sodium chloride and ammonium chloride do not. 

Solutions of the hydrated salt have exactly the same prop- 
erties as those of solutions of the anhydrous salt. 

98. Sodium Hydrogen Sulfate: NaHS0 4 . — If we exactly 
neutralize a definite quantity of sulfuric acid with a solution of 
sodium hydroxide, noting the volume of the latter used, and again 
add to a second portion of sulfuric acid, exactly equal to the first, 
exactly half as much sodium hydroxide solution as that used 
in the first case, we find that the first solution yields when 
evaporated pure sodium sulfate, Na 2 S0 4 ; while the second gives 
crystals having a different shape and appearance, and different 
chemical properties. Analysis shows that the composition of 
these crystals is represented by the formula NaHS0 4 . The 
substance is called sodium hydrogen sulfate. The equation for 
the reaction in the second case is 

H 2 S0 4 +NaOH-> NaHS0 4 +H 2 0. 

99. The Law of Definite Composition Again.— We may now 

consider one of the most important and fundamental of all 
chemical questions, namely, whether the proportions of the 
elementary constituents of a substance are dependent upon the 
proportions which we take to the substances from which we form 
the substance in question. For example, we may inquire 
whether we could get a sulfate of sodium with a somewhat 
larger or smaller percentage of sodium if we had used, in the 
preceding experiment, other proportions of acid and base. 
Experiment will show, however, that if we had added a little 
more or less sodium hydroxide we would still have been able to 
obtain much XaHS0 4 , but that in such cases there would also 
be some Na 2 S0 4 formed or a little free sulfuric acid left after all 



Acids, Bases, and Salts — / 63 

the NaHS0 4 had been separated from the water. Facts like 
these which are met with on every hand give a special significance 
to the Law of Definite Composition. 

100. Acid Properties of Sodium Hydrogen Sulfate.— We see 
that sodium sulfate, Na 2 S0 4 , contains exactly twice the weight of 
sodium for a given weight of sulfur and oxygen as does sodium 
hydrogen sulfate, NaHS0 4 . Moreover, we have become 
acquainted with the important fact that sulfuric acid can form 
two sorts of sodium salts. If we dissolve crystals of sodium hydro- 
gen sulfate in water, we find that the dilute solution has a sour 
taste and it turns litmus red, for which reasons we should be 
inclined to say that it has acid properties. In accord with this 
view, we find that the solution will readily neutralize a solution 
of sodium hydroxide, giving sodium sulfate and water, thus: 

NaHS0 4 -f NaOH -> Na 2 S0 4 +H 2 0. 

1 01. Ammonium Sulfate and Ammonium Hydrogen Sul- 
fate. — If we completely neutralize sulfuric acid with a solution 
of ammonium hydroxide, we obtain a salt called ammonium sul- 
fate (NH 4 ) 2 S0 4 , thus: 

H 2 S0 4 +2NH 4 OH-> (NH 4 ) 2 S0 4 +2H 2 0; 

while with half the proportion of ammonium hydroxide we obtain 
ammonium hydrogen sulfate, thus : 

H 2 S0 4 +NH 4 OH -> NH 4 HS0 4 +H 2 0. 

102. Monobasic and Dibasic Acids: Acid Salts and Neutral 
Salts. — Hydrochloric acid reacts with sodium hydroxide only in 
one proportion, thus: 

HCl+NaOH -» NaCl+H 2 0, 

for which reason we call it a monobasic acid; but since one 
formula weight of sulfuric can unite with a maximum of two 
formula weights of sodium hydroxide we call sulfuric acid a 
dibasic acid. Salts in which but half the maximum quantity 
of base has been neutralized are usually called acid salts, because 
they still have acid properties. Thus we frequently speak of 
sodium acid sulphate, meaning NaHS0 4 . Chemists know many 



6 4 



Introduction to General Chemistry 



other dibasic acids, all of which also can form acid salts as well as 
neutral salts, as salts like Na 2 S0 4 are called. 

103. Making Hydrochloric Acid from Common Salt. — If we 
place in a flask (Fig. 23) 58 g. of dry common salt and 100 g. of 
sulfuric acid, to which 30 g. of water have been added, and warm 
the mixture, a change occurs with the production of a colorless 
gas which dissolves in water very readily, giving a solution which 
we can easily recognize as hydrochloric acid.' After the action 
of the sulfuric acid on the salt is complete, a white solid is left in 
the flask, which may easily be dissolved in water. By evaporat- 
ing part of the water, and letting the solution stand a while, we 

may obtain colorless, transparent 

crystals of sodium hydrogen 
sulfate. The following equation 
represents the reaction: 

NaCl+H 2 S0 4 -» NaHS0 4 +HCl. 



We have to deal here with a new 

sort of chemical change — one in 

which an acid acts upon a salt of 
another acid to give a salt of the 
first acid and to produce the acid corresponding to the first salt. 
This is a very important kind of chemical reaction, which we 
shall frequently make use of, since by its means we may make 
acids from their salts. 

104. Making Nitric Acid from Chile Saltpeter. — We shall now 
use the method just described for the preparation of a new acid 
from a white, crystalline substance called Chile saltpeter, which 
is found in large quantities as a mineral substance in the desert 
region of Chile. 

If we place 85 g. of Chile saltpeter in a retort (Fig. 24), add 
100 g. of sulfuric acid, mixed with 30 c.c. of water, and then heat 
the mixture gently, a yellow-colored liquid may be collected in a 
cooled flask. This yellow liquid gives off a brown gas and 
becomes colorless when boiled a few minutes. Its analysis shows 
its formula to be HN0 3 and it is called nitric acid. It is a color- 
less liquid which may be boiled and distilled in glass vessels. 




Fig. 23 



Acids, Bases, and Salts — / 



65 



Pure or concentrated nitric acid is even more dangerous than 
sulfuric acid, causing serious burns and destroying clothing, 
and must be handled with greatest care. It will mix with water 
in all proportions, giving a solution which, when very dilute, has 
a sour taste and turns litmus red. 

When nitric acid is mixed with sodium hydroxide solution the 
latter is neutralized, a salt of the composition NaN0 3 and water 
being the only products, as represented by the equation 

HN0 3 +NaOH -> NaN0 3 +H 2 0. 

The salt, which is called sodium nitrate, is found to be identical 

with purified Chile saltpeter. The 

action of sulfuric acid on saltpeter 

leaves in the retort a white solid 

which closely resembles that left 

when salt is heated with sulfuric acid, 

and, in fact, the residue is easily found 

to be the same substance, sodium 

hydrogen sulfate, NaHS0 4 . The 

equation for the reaction is therefore 




Fig. 24 



NaN0 3 +H 2 S0 4 -> NaHS0 4 +HN0 3 . 

105. The Action of Nitric Acid on Ammonium Hydroxide. — 

We may now propose a question to be answered, not after direct 

experiment, but as a result of the general knowledge we have 

gained regarding the behavior of the acids and bases already 

studied. I«t is: What would be the result of mixing nitric acid 

and ammonium hydroxide? We recall that hydrochloric acid 

and sodium hydroxide, a base, give sodium chloride and water, 

thus: 

HCl+NaOH-> NaCl+H 2 0; 

that the same acid gives with ammonium hydroxide, also a base, 
ammonium chloride and water, thus: 

HCl+NH 4 OH ^ NH 4 C1+H 2 0. 

Furthermore, we have just seen (104) that nitric acid and sodium 
hydroxide give sodium nitrate and water, thus : 
HN0 3 +NaOH -> NaN0 3 +H 2 0, 



66 Introduction to General Chemistry 

and we would certainly expect that nitric acid and ammonium 
hydroxide would behave analogously and give ammonium nitrate 
and water, thus : 

HN0 3 +NH 4 OH -> NH 4 N0 3 +H 2 0. 

Now this is precisely what takes place when we test our prediction 
by experiment. We seem, therefore, to have discovered the 
secret of the way in which acids and bases act toward each other. 
It may be summed up in the statement, An acid and a base 
neutralize each other, forming a salt and water. 

106. A New Base: Caustic Potash or Potassium Hydroxide. 
— Let us now take up the study of a new base, caustic potash, 
which closely resembles caustic soda (sodium hydroxide). It 
will be remembered that the metal sodium reacts violently with 
water, giving sodium hydroxide and hydrogen gas, thus: 

2 Na+ 2 H 2 -> 2 NaOH+H 2 (88) 

Now, chemists know another metallic element, potassium, which 
closely resembles sodium. Like sodium, it is a silver-white 
metal, soft enough to be cut easily with a knife and tarnishing 
very rapidly in the air. For a reason that we shall soon learn it 
is kept covered with oil in a carefully stoppered bottle. If we 
throw a small bit of potassium into a beaker of water, it bursts 
into a flame of lavender color, spinning and darting to and fro 
on the surface of the water and completely disappearing in a few 
moments. Examination of the water shows that it will turn 
litmus blue, that it has a " soapy" taste, like a very dilute solu- 
tion of sodium hydroxide, and that a white solid is left when the 
solution is evaporated to dryness. This solid is found by suitable 
methods of analysis to contain the elements potassium, oxygen, 
and hydrogen in the proportion represented by the formula 
KOH, and is called potassium hydroxide. 

If instead of throwing the bit of potassium on the surface of 
the water we bring it under the mouth of an inverted cylinder 
filled with water, with the mouth immersed in a vessel of water, 
the potassium rises to the top of the water in the cylinder, pro- 
ducing a gas which displaces the water in the cylinder, but does 



Acids, Bases, and Salts — / 67 

not take fire. The gas is easily identified as hydrogen, while the 
water contains dissolved potassium hydroxide as before. The 
equation for the reaction in the cylinder is 

2 K+2H 2 0->2KOH+H 2 . 

When the action takes place in the open beaker, the heat pro- 
duced sets fire to the hydrogen, which burns, together with a small 
portion of the potassium. 

107. Potassium Salts. — On account of the behavior of a solu- 
tion of potassium hydroxide toward litmus and also because of its 
"soapy" feel and taste, we should conclude that it is a base and 
if so that it should form salts with acids. We might even venture 
to predict the formulae of the salts it would be expected to form 
with hydrochloric, sulfuric, and nitric acids, and to write the 
equations as follows: 

HCl+KOH-> KC1+H 2 
H 2 S0 4 + 2KOH -> K 2 S0 4 + 2 H 2 
H 2 S0 4 +KOH -> KHS0 4 +H 2 
HNO3+KOH -> KN0 3 +H 2 0. 

And in every case these predictions would be found by experiment 
to be correct! The potassium salts so formed are all white 
crystalline solids and are all soluble in water. All except potas- 
sium hydrogen sulfate give solutions which are neutral to litmus, 
while this salt has acid properties like those of sodium hydrogen 
sulfate. 



CHAPTER VIII 
WATER AND SOLUTIONS 

1 08. Water. — We have already learned that pure water is 
readily obtained by the distillation of natural waters (23), and 
that it is a compound of hydrogen and oxygen, the composition 
of which is represented by the formula H 2 (70) . In describing 
a substance we shall often mention its physical and chemical 
properties. The properties of a substance embrace: the state 
(whether solid, liquid, or gaseous); crystalline form, if solid; 
specific gravity or density ; color; odor; taste; conductivity for 
heat and electricity; boiling-point; freezing-point, etc. The 
chemical properties of a substance are those which it exhibits in 
its typical chemical reactions. 

109. The Physical Properties of Water: Color. — We know 
that according to the temperature water can exist as solid, liquid, 
or gas. The color of liquid water is a very faint blue; so faint, in 
fact, that it cannot be noticed in a glass of water, but is obvious 
in a white bathtub full of clear water. The color of large bodies 
of clear water is usually blue, but it may be of any other shade 
if dissolved or suspended impurities (mud) are present. The 
yellow color of the waters of many rivers is due to suspended 
clay; such water is not clear, but muddy or turbid. Streams and 
lakes in hemlock forests often contain perfectly clear water 
having the color of tea, due to coloring-matter dissolved from 
the hemlock. The clear green color of some waters is usually 
the result of the blending of the natural blue color of the water 
with the yellow light reflected from the sand beneath. 

no. Specific Gravity or Density. — At the temperature of 
4° C, 1 c.c. of water weighs 1 g. Since the specific gravity or 
density of any substance may be defined as the weight of 1 c.c, it 
follows that water has a specific gravity of 1 . 000 at 4 C. Or, we 
may say that the specific gravity or density of a substance is 
found by dividing its weight by the weight of an equal volume of 
water. Water has its greatest density at 4 ; if a given volume 

68 



Water and Solutions 



69 






of water at 4 is either heated or cooled, it expands and therefore 
decreases in density. 

in. Specific Heat. — The quantity of heat required to raise 
the temperature of 1 g. of water i° C. is by definition called one 
calorie. Water is said to have a specific heat of one or unity. 
The specific heat of any substance is the quantity of heat in calories 
required to raise the temperature of one gram of it one degree. 
Nearly all substances have specific heats less than unity. 

112. Vapor Pressure. — Water contained in an open vessel 
evaporates at all temperatures, but the more rapidly in propor- 
tion as the temperature is higher, other things being equal. If 
water evaporates into an evacuated space the 
pressure within the space increases to a value 
which is dependent only upon the tempera- 
ture, being greater in proportion as the 
temperature is higher. The pressure so pro- 
duced is called the vapor pressure of water; 
it may easily be demonstrated by means of a 
barometer tube filled with mercury. If we 
prepare two such tubes (Fig. 25) and intro- 
duce a few drops of water into one by means 
of a suitably shaped glass tube, the water 
will rise until it floats on the surface of the 
mercury. At the same time the level of the mercury will 
fall 2 or 3 cm., showing that a pressure has been produced 
above the mercury in the space which has been a vacuum. 
If the tube into which the water is introduced has a glass jacket 
into which warm water can be poured, it will be found that 
the higher the temperature is, the higher the vapor pressure will be. 
If we should raise the temperature to ioo°, the level of the 
mercury in the barometer tube would sink to that of the surface 
of the mercury in the dish in which the tube stands, thus show- 
ing that the vapor pressure at ioo° is equal to the pressure of the 
atmosphere. Table VII shows the vapor pressure of water at 
various temperatures between o° and ioo°. 

When the atmospheric pressure is 760 mm., water boils at 
ioo°. Now, we see from the table that at ioo° the vapor 




Fig. 25 



7° 



Introduction to General Chemistry 



pressure is 760 mm., therefore the boiling-point is that temperature 
at which the vapor pressure becomes just equal to the normal atmos- 
pheric pressure, 760 mm., which is the average pressure at sea- 



TABLE VII 



Temperature 


Pressure 


Temperature 


Pressure 


o° 

10 

20 

30 

40 

5o 


4 . 6 mm. 
9.2 

17-4 
31.6 

55-0 
92. 2 


6o° 

70 

80 

9° 

99 

100 


149. 2 mm. 
233-8 

355-5 
526.0 
733-2 
760.0 



level. At higher altitudes, at which the atmospheric pressure 
is less than 760 mm., water boils at temperatures lower than 
ioo°. Thus if the pressure is 733.2mm., the boiling-point is 
99 . Since the atmospheric pressure at a given place is variable 
through a range of 20 mm. or more, the boiling-point at this 
place is not constant, but varies with the rise and fall of the 
barometer. 

113. Correction of the Volume of a Gas for Vapor Pressure. — 
Gases like hydrogen and oxygen, which are not very soluble in 
water, are often measured in tubes in which the 
gases are confined by means of water. Such gases 
always contain water vapor, and part of the total 
pressure exerted by the gas is due to the vapor 
pressure of the water. The part of the pressure 
(partial pressure) exerted by the gas itself is found 
by subtracting from the total pressure the vapor 
pressure of the water. For example, suppose that 
some hydrogen is collected over water in a grad- 
uated glass tube (Fig. 26). If the position of the 
tube is adjusted so that the level of the water is the same inside 
the tube as outside, the total pressure within must be exactly 
equal to the atmospheric pressure, as shown by the barometer. 
Suppose that the barometric pressure is 748 . 6 mm. and the 
temperature 20 . Table VII shows that at 20 the vapor 
pressure of water is 17.4 mm., therefore the pressure due to the 



Fig. 26 



Water and Solutions 71 

hydrogen is 748.6—17.4 = 731.2 mm. If the observed volume 
was 30 c.c, the volume, V, at standard conditions would be 

F = 3°X|3i_l>073 = 26 

760X293 

114. Vapor Pressure of Liquids and Solids in General. — 

Liquids in general readily pass into the form of vapor, and just 
as in the case of water, a given pure liquid has, at each tempera- 
ture, a definite vapor pressure; but the vapor pressure of one 
liquid — say alcohol — is not in general the same at a given 
temperature as that of another liquid — say water. In every 
case, however, the boiling-point of the liquid is that temperature at 
which its vapor pressure equals j6o mm. Many solids, for 
example, camphor and naphthalene (moth-balls), have appre- 
ciable vapor pressures at room temperature; but the vapor 
pressures of most solids at such temperature are too small to be 
noticeable. 

115. Latent Heat of Evaporation. — If it is true that water 
boils at ioo° because at this temperature the vapor pressure of 
water just equals the normal atmospheric pressure, it may be 
asked why the whole of the water does not change at once into 
steam as soon as its temperature is raised to ioo°. We know, of 
course, that this does not occur, and, further, that the rapidity 
with which water boils away is greater, the greater the amount 
of heat applied. The explanation is found in the fact that it 
requires a large amount of heat to change water at ioo° into steam 
at the same temperature. In fact, 540 calories of heat are required 
for the conversion of 1 g. of water at ioo° into steam. The heat 
so used up does not raise the temperature of the substance. It is 
consumed in changing the liquid water into the gaseous state; 
it is said to become latent, and in consequence we say that the 
latent heat of evaporation of water is f&o- calories. Every pure 
liquid has a latent heat of evaporation. This differs from one 
substance to another. 

116. Use of Steam for Heating.— When steam cools to ioo° it 
begins to condense to liquid water, and for every gram of steam 
that condenses 540 calories of heat are given out. The heat 



72 Introduction to General Chemistry 

so given out may be considered to be that which became latent 
when the water was, by being heated, converted into steam. It 
is on account of the latent heat given out upon condensation that 
steam is so effective in the heating of buildings: every gram of 
steam that condenses in the radiator liberates 340 calories of 
heat. Of course, the further cooling of the water in the radiator 
gives out some additional heat. 

117. Burns Produced by Steam. — It is a well-known fact 
that serious burns result when steam comes in contact with the 
skin. At first thought, this result seems to be out of harmony 
with the fact that air at ioo° can be borne by the hand without 
discomfort. The explanation of this difference is found in the 
fact that gases (including the vapors of boiling liquids) are very 
poor conductors of heat as compared with liquids. Steam at ioo° 
partly condenses on striking the skin and wets it with a layer 
of boiling-hot water, which is a good conductor of heat . . Further- 
more, since 540 calories of heat are given out by every gram of 
steam condensed to water, the latter is kept at ioo° as long as 
steam is present. On the other hand, air is so poor a conductor 
of heat that the skin is not burned by a brief exposure to it at ioo°. 

118. Latent Heat of Fusion of Ice. — Ice melts at o°; but all 
of a given mass of ice does not melt immediately when its 
temperature is raised to zero. Just as heat is required to change 
liquid water into vapor, so also heat is needed to change ice at 
zero into water at the same temperature. The heat so absorbed 
is called the latent heat of fusion of ice. It requires 79 calories 
to melt 1 g. of ice; therefore the latent heat of fusion of ice is 
79 calories. Every solid has a definite and characteristic latent 
heat of fusion. 

119. The Density of Ice. — The density or specific gravity of 
ice is 0.917. It is for this reason that ice floats on water. The 
expansion which occurs when water freezes exerts very great 
pressure, illustrations of which are often seen in the bursting of 
water pipes and other vessels when water freezes in them. Not 
all liquids expand upon freezing; in many cases contraction 
occurs, thereby giving rise to solids which sink in the correspond- 
ing liquids. 



Water and Solutions 73 

120. Solutions and Suspensions. — The mixture which results 
upon dissolving salt in water is called a solution of salt in water. 
The terms " dissolve" and "solution" are used in chemistry with 
definite meanings. If, upon mixing a solid with a liquid, the 
former partly or wholly disappears and the resulting liquid is 
still clear and transparent and not cloudy or muddy, and if, 
moreover, upon allowing the liquid to evaporate we regain the 
unchanged solid substance, we say that the solid had dissolved 
in the liquid to form a solution. Either or both of the substances 
may be colored and still a clear (although colored) solution may 
result. The liquid in which a substance is dissolved is called the 
solvent. 

If we stir up some common clay with water, much of the clay 
fails to settle out of the water at once, and we get a cloudy or 
muddy fluid, like the water of a muddy river. In this case we 
do not say that the clay has dissolved in the water or that we have 
a true solution of the clay. We say that the clay is suspended 
in the water, and call the muddy water a suspension. Clay 
suspended in water will settle out very slowly and finally leave 
clear water above a layer of mud. 

121. The Concentration of Solutions. — A solution containing 
a small proportion of a dissolved substance is said to be dilute, 
while one containing a large proportion is called concentrated. 
We dilute a concentrated solution by adding solvent to it, and 
concentrate a dilute solution by evaporating the solvent. We use 
the term concentration in discussing the relative amount of 
dissolved substances in a solution. 

122. Solubility of Substances: Saturated Solutions. — It is 
easy to discover that the amount of a substance which will dis- 
solve in a given amount of water, say 100 c.c, depends upon the 
nature of the substance and upon the temperature. If we mix 
some common salt with about double its weight of water and 
stir or shake the mixture a sufficient length of time (usually one 
to two hours), keeping the temperature constant all the while, 
and then, after allowing any suspended crystals to settle, draw 
off a portion of the clear solution, weigh it, and evaporate the 
water, we get the salt dissolved in the portion of the solution 



74 Introduction to General Chemistry 

taken. By weighing the salt we can readily find the weight of 
salt dissolved in a given weight of water at the temperature at 
which the experiment was made. We find in this way that 
ioo g. of water at 25 dissolves 37.6 g. of salt. 

To make such a solubility determination we must observe 
several precautions: First, the amount of solid substance must 
be considerably greater than the amount of water taken will dis- 
solve; secondly, the shaking must be continued as long as more 
substance dissolves — this is easily ascertained by prolonging the 
shaking and making additional determinations of the concentra- 
tion of the solution; thirdly, the temperature must be kept 
constant. 

A solution which at a fixed temperature will dissolve no more 
of a given substance is called a saturated solution. When we 
speak of the solubility of a substance we mean the amount of 
substance dissolved in a given amount of water in the case of 
the saturated solution. The following brief table gives the 
solubilities in water at 25 of several salts. 

TABLE VIII 

Grams of Substance in ioo g. of Water at 25 



NaCl 


37 g- 

27 
92 


KC1 


34 g. 
12 

37 


Na 2 S0 4 ■ ioH 2 
NaN0 3 


K 2 S0 4 

KNO3 



123. Supersaturated Solutions. — At 25 100 g. of water will 
dissolve 27 g. of sodium sulfate decahydrate, Na 2 S0 4 * ioH 2 0, 
while at 30 the same amount of water will dissolve 40 g. of the 
salt. If we make a saturated solution of the salt at 30 , having 
an excess of crystals of the salt present, and then cool the whole 
to 25°, and keep it at 25 , stirring or shaking it for an hour or two, 
more solid is deposited and there results a solution which contains 
just the same weight of the salt in 100 g. of water as a saturated solu- 
tion at 2 5 , namely, 27 g. 

A slight change in the procedure gives a very different result 
and brings to light a new phenomenon. If the solution of sodium 
sulfate which is saturated at 30 is freed from every particle of the 



Water and Solutions 75 

solid crystalline substance and then allowed to cool to 25 or even 
lower, without being stirred or shaken, it remains perfectly clear and 
does not deposit any crystals. Such a solution contains at 25 
much more sodium sulfate than a saturated solution prepared 
at 2 5 in the manner described in the preceding paragraph. 
This more concentrated solution is called a supersaturated 
solution. If we now drop into the supersaturated solution a 
crystal of sodium sulfate (and for this purpose an almost in- 
visible fragment of the crystalline dust will be sufficient), the 
formation of crystals will begin at once and proceed until the 
amount of dissolved substance per 100 g. of water is reduced 
exactly to that of a saturated solution at the existing tempera- 
ture. 

Experience has shown that a supersaturated solution can 
only be obtained in the complete absence of the solid substance, 
and that a supersaturated solution begins to deposit its excess 
of dissolved substance when a crystal of this same substance is 
brought into the solution. The deposition of crystals by a 
supersaturated solution can also often be started by shaking 
or stirring the solution or by adding a crystal of another sub- 
stance having the same crystalline form. 

Not all substances form supersaturated solutions equally 
readily. The presence of impurities favors supersaturation. 
Syrups, preserves, and honey are often supersaturated with 
respect to the sugar dissolved in the water present. When such 
solutions "turn to sugar," this is only the crystallization of the 
excess of sugar above that required to make a saturated solution. 

124. Solubility of Liquids in Liquids. — It is proverbial that 
"oil and water will not mix." On the other hand, some pairs of 
liquids will mix completely in all proportions; examples of such 
combinations are water and alcohol and water and sulfuric acid. 
We know other pairs of liquids that will not dissolve one another 
in all proportions, but that will dissolve one another partially. 
Water and ether belong to this class; 100 c.c. of water will 
dissolve 8 c.c. of ether, and 100 c.c. of ether will dissolve 3 c.c. 
of water. If we pour ether into water, we find that the former 
floats on the surface of the latter. If equal volumes of ether and 



76 Introduction to General Chemistry 

water are thoroughly shaken together, the former soon separates 
from the latter, and two distinct layers result as before. If, now, 
we examine each layer, we find that the water contains some 
dissolved ether and the ether some dissolved water. This is a 
case of partial miscibility. 

125. Solubility of Gases in Liquids. — We have already 
learned that hydrogen chloride (44) and ammonia (51) are both 
very soluble in water. At o° water dissolves 550 times its own 
volume of the first gas and 1,150 times its volume of the second. 
No gas which we have studied is completely insoluble in water; 
for example, 100 c.c. of water dissolves 2. 1 c.c. of hydrogen and 
4.8 c.c. of oxygen. Fishes depend for their existence upon the 
oxygen dissolved in water; by means of their gills they take 
from the water the oxygen they require. 

126. Henry's Law. — The solubility of all gases decreases with 
rise of temperature. At a fixed temperature the weight of gas 
dissolved by a given volume of water of other liquid is dependent 
upon the pressure of the gas and is, in general, directly propor- 
tional to the pressure. This statement is known as Henry's 
Law. The law does not apply to very soluble gases, like am- 
monia, dissolving in water — probably because chemical union 
occurs, since we know that XH 4 OH is formed in this case (91). 

127. Heat of Solution. — If we shake some potassium nitrate 
or ammonium chloride, or indeed any one of many salts, with 
water, we find that as the substance dissolves the solution becomes 
appreciably colder. This indicates that heat is required to change 
the solid into the dissolved state. This phenomenon is analogous 
to that met with when a solid, like ice. melts. It requires 79 
calories to melt 1 g. of ice, while 115 calories are absorbed when 
1 g. of potassium nitrate dissolves. That is, we must supply 
115 calories to 1 g. of the salt, and sufficient water, in order to 
prevent a fall of temperature when solution takes place. The 
heat so required is called heat of solution. 

When any substance whatever melts, heat is required, or is 
absorbed, and we might expect, similarly, that heat will always 
be absorbed when a substance dissolves; but this is not the case. 
Many substances, upon dissolving, give out heat. In the case of a 



Water and Solutions 77 

few substances the absorption or evolution of heat upon dissolv- 
ing is very small. Common salt dissolves in water with very 
small heat absorption. 

128. Boiling-Point of Solutions. — It is very easy to show that 
a solution of a solid substance, like salt or sugar in water, boils 
at a higher temperature than pure water. This is an invariable 
rule for solutions of substances which are not readily volatile 
at the boiling-point of water. Now, we have in the first part of 
this chapter (112) considered the relationship between boiling- 
points and vapor pressures, and it will easily be understood that 
a solution will boil at the temperature at which the pressure of 
its vapor is equal to the atmospheric pressure. 

129. The Lowering of the Vapor Pressure by Dissolved Sub- 
stances. — If a solution must be heated above ioo° to raise its 
vapor pressure to that which water has at ioo°, it is clear that at 
this latter temperature the solution has a lower vapor pressure than 
pure water. It is also a fact that at every lower temperature the 
vapor pressure of a solution of an involatile substance is less than 
that of the pure solvent at the same temperature. This is a 
very important universal law. The law applies to solutions 
formed from all kinds of solvents. 

130. Deliquescence. — In the case of a very soluble substance, 
like caustic soda, the vapor pressure of the saturated solution 
may be so small that it is below the partial pressure exerted by 
the vapor usually present in the air. If such a solution is 
exposed to the air, water vapor from the air will condense in it 
until the solution has become so dilute that its vapor pressure 
is just equal to the partial pressure of the water vapor in the air. 
Moreover, if such a very soluble substance is exposed to air con- 
taining moisture, water will condense on the solid, thus convert- 
ing it slowly, first into a saturated solution, and finally into a 
dilute solution. This action is called deliquescence. We say 
caustic soda is a deliquescent substance. A little thought will 
lead to the conclusion that deliquescence is the result of two con- 
current conditions; first, the possibility of the formation, by a 
substance, of a saturated solution which has a very small vapor 
pressure as compared with pure water — a condition usually 



78 Introduction to General Chemistry 

accompanying great solubility; and, secondly, the presence in the 
air of a sufficiently great water-vapor content. No substance 
is deliquescent in a perfectly dry atmosphere, while every 
soluble substance exhibits this property in air saturated with 
water vapor. Deliquescence is, therefore, not a fixed property of 
a substance. Thus common salt is usually decidedly deliquescent 
at the seashore, where the air contains much water vapor; but 
it never shows this property in a desert region. 

In several experiments we have used caustic soda or calcium 
chloride to dry air or other gases or to absorb water vapor formed 
in the burning of hydrogen (39, 50). These drying agents are 
among the most deliquescent substances known. 

131. Efflorescence. — In paragraph 94 the peculiar behavior 
of sodium sulfate decahydrate, Na 2 S0 4 , ioH 2 0, when exposed 
to the open air was described. We are now in a position to 
understand more about this spontaneous loss of water. If a 
crystal of the hydrate is floated on the surface of mercury in a 
vaccum tube like one of those shown in figure 25, the mercury 
level is depressed more than can be accounted for by the weight 
of the crystal. Apparently the latter is giving off water vapor 
and attempting to establish a saturation pressure. This pres- 
sure is called the vapor pressure of the hydrate. As a matter of 
fact all hydrates show this same behavior, with the difference 
that each has its own characteristic vapor pressure at a given 
temperature. With increased temperature the vapor pressure 
rises. If a hydrate is exposed to air in which the partial pressure 
of water vapor is less than the vapor pressure of this substance, 
the latter will give off water to the air just as a water surface 
does to air in which the partial pressure of water vapor is below 
the saturation value for water. Along with the loss of water, 
the crystals of the decomposing hydrate crumble to a powder. 
This process is called efflorescence. It is obvious that whether 
or not a given hydrate effloresces depends not only upon its own 
vapor pressure but upon the moisture content of the air surround- 
ing it. 

132. Effect of Temperature on Solubility. — The solubility of 
a substance, that is, the amount of the substance which dissolves 



Water and Solutions 



79 



(to form a saturated solution) in a given amount of water, is 
dependent upon the temperature. Most substances are more 
soluble at a higher than at a lower temperature; but this is not 
always the case, as the solubility of some substances decreases 
with rise of temperature. In fact, gases are always less soluble 
at a higher temperature. 




■*0 J50 bO 



Fig. 27 



The change of solubility with change of temperature can 
most easily be expressed graphically, that is, by means of 
so-called solubility curves. The accompanying diagram (Fig. 2 7) 
illustrates the method and gives the curves for water solutions of 
several substances. 

133. Effect of Crystalline Form on Solubility.— Sodium sul- 
fate has the formula Na 2 S0 4 . By the action of water we may 



80 Introduction to General Chemistry 

readily obtain the hydrate Na 2 S0 4 *ioH 2 (96), which can easily 
be recrystallized from water, as described under " Supersaturated 
Solutions" in this chapter. We see that the solubility curve for 
Na 2 S0 4 * ioH 2 rises rapidly until a temperature of 33 is reached. 
At this temperature the crystals melt and at the same time 
decompose into Na 2 S0 4 and H 2 0, thus: 

Na 2 S0 4 • ioH 2 -> Na 2 S0 4 + ioH 2 0. 

Above 33 we have the solubility curve of anhydrous Na 2 S0 4 
which is a different chemical substance from its hydrate. Thus 
we see that there are for the anhydrous salt and its hydrate two 
distinct solubility curves, and that these intersect at a point 
for which the temperature is that at which the hydrate changes 
into the anhydrous substance. This is a typical case. Each 
hydrate of a substance has its own solubility curve; but these always 
intersect at the point corresponding to the temperature at which one 
substance changes into the other. The difference in solubility is 
due to the fact that each has its own characteristic crystalline 
form. 

134. Heat of Solution and Changes of Solubility with 
Temperature. — A question which will now very naturally occur 
to the student is: Why should the solubility of various sub- 
stances change with temperature in different ways? Although a 
complete and satisfying answer cannot be given to this question, 
it is possible to find a connection between the shape of the solu- 
bility curve of a substance and another fundamental property. 
It will be recalled that potassium nitrate absorbs much heat 
upon dissolving in water, and we notice that its solubility curve 
rises rapidly with temperature. Sodium chloride dissolves with 
but slight absorption of heat, and its curve is nearly horizontal. 
Finally, when it is known that anhydrous sodium sulfate, 
Na 2 S0 4 , dissolves at temperatures above 33 with production of 
heat, and that its curve falls with rising temperature, the general 
law becomes apparent. These are typical cases. If any sub- 
stance dissolves with absorption of heat, its solubility curve 
rises with rise of temperature. If it dissolves with evolution of 
heat, then the curve falls with rise of temperature. The frac- 



Water and Solutions 81 

» 
tional change of solubility with rise of i° of temperature is in 

general proportional to the heat of solution. In every case that 
change of solubility which will absorb heat will take place when the 
temperature is raised. This will involve a decrease of solubility 
with rise of temperature, in the case of a substance like Na 2 S0 4 , 
above 33 , since, if heat is evolved when the substance dissolves, 
heat is absorbed in equal amount when the same weight of the 
substance crystallizes out of a solution. 

In some cases where heat is evolved when a substance is 
dissolved, the observed heat is the result of the union of the solid 
with water to form a hydrate, which may dissolve with a small 
absorption of heat. In such cases the solubility of the hydrate 
increases with rise of temperature in strict accord with the law. 
For example when anhydrous calcium chloride is dissolved in 
water the mixture gets very hot. The saturated solution 
deposits crystals of CaCl 2 , 6H 2 on cooling. This hydrate dis- 
solves in water with absorption of heat and its solubility increases 
with a rise in temperature. The heat given out on dissolving the 
anhydrous salt is the excess of the heat produced in the reaction 

CaCl 2 +6H 2 = CaCl 2 ,6H 2 

above the heat absorbed in the dissolving of the hydrate CaCl 2 , 
6H 2 0. 

135- Two Apparent Kinds of Solubility. — In cases of ordinary 
solubility, evaporation of the water leads to the recovery un- 
changed of the substance originally dissolved. In other cases, 
evaporation of the solution obtained by the apparent dissolving 
of a substance leaves an entirely different substance. For 
example, if we throw a piece of sodium on water the former soon 
disappears and a solution results (40, 88) . We might be inclined 
to say that the sodium has dissolved in the water; but there is 
another way of looking at the matter. We know that in this 
case a chemical change has occurred, as represented by the 
equation 

2 Na+2H 2 0^ 2 NaOH+H 2 . 

Furthermore, we know that by evaporation of the solution we 
get sodium hydroxide and not sodium; for this reason it seems 



82 Introduction to General Chemistry 

more logical to say that sodium and water react to give sodium 
hydroxide, which then dissolves in water, than to say that 
sodium itself is soluble in water. In fact, we know nothing 
about the solubility of sodium in water, since the two react as 
soon as they are brought into contact. We know a very great 
number of cases analogous to this one, and in all of them we 
recognize that we have to deal with chemical changes which 
give rise to soluble products. 

136. Normal Solutions. — In the neutralization of hydro- 
chloric acid by sodium hydroxide, which takes place according 

to the equation 

HCl+NaOH -> NaCl+H 2 0, (89) 

one formula weight of the acid (36. 5 g.) requires one formula 
weight of the base (40 g.). If we make a solution of the acid of 
such concentration that 1 liter contains 36. 5 g. of hydrogen 
chloride, and also make a solution of the base containing 40 g. 
of sodium hydroxide per liter, then upon mixing the liter of the 
acid solution with the liter of the basic solution exact neutraliza- 
tion will take place. It follows, of course, that, to neutralize a 
given volume of such an acid solution, exactly the same volume 
of the basic solution will be required. We call such solutions 
normal solutions. 

If we wish to make a solution of nitric acid of such concentra- 
tion that 1 liter of it will exactly neutralize 1 liter of normal 
sodium hydroxide, we see, in accord with the equation 

HN0 3 +NaOH -> NaN0 3 +H 2 0, (104) 

that one formula weight of HN0 3 must be contained in 1 liter 
of the solution. This gives a normal solution of nitric acid. 

Now the case is a little different if a normal solution of sul- 
furic acid is to be made, since in this case we have 

H 2 S0 4 + 2 NaOH -> Na 2 S0 4 + 2H 2 0. (94) 

We see that one formula weight of sulfuric acid neutralizes two 
formula weights of sodium hydroxide, so that to neutralize 
1 liter of normal sodium hydroxide, which contains but one 
formula weight of the base, only one-half a formula weight 



Water and Solutions 



83 



«5CC 

AJ«C 



Fig. 28 



(J of 98 g. or 49 g.) of sulfuric acid is required. Therefore if we 
dissolve 49 g. of the acid in sufficient water to make a liter of 
solution, this liter of acid solution will just neutralize 1 liter of 
normal sodium hydroxide. We call the sulfuric 
acid solution so made also a normal solution. 

A normal solution of potassium hydroxide, 
KOH, would contain one formula weight (56 g.) 
per liter (106). A normal solution of any acid 
always neutralizes an exactly equal volume of a 
normal solution of any base. The term " normal" 
is usually abbreviated N, so that for a normal 
solution by hydrochloric acid we write N HC1. 
Normal solutions are of great importance in prac- 
tical work. Suppose we wish to know the concen- 
tration of a given solution of sodium hydroxide. 
We take, with a pipette (Fig. 28), a carefully 
measured volume, say 20 c.c, add to it sufficient 
litmus solution to produce a pale blue color, and then from a 
measuring tube, called a burette (Fig. 29), run in a normal solu- 
tion of hydrochloric or other acid until the color just changes 
from blue to red. A little practice enables one to find, to within 

one drop or less, the volume of acid 
required. Let us say 42 c.c. of 
N HO was required for the 20 c.c. 
of NaOH solution of unknown con- 
centration. Our problem is to find 
the weight of sodium hydroxide in 
the 20 c.c. of solution taken. Now, 
42 c.c. of N acid will neutralize 
42 c.c. of N sodium hydroxide, of 
which 1 liter (=1,000 c.c.) contains 
40 g. of sodium hydroxide. There- 
fore the weight of sodium hydroxide in the 20 c.c. taken = o . 04.2 X 
40 g. = 1 . 68 g. We also see that the sodium hydroxide solution 
is 42/20 = 2. 1 times as concentrated as a normal solution of this 
base. We express its concentration by saying that it is 2.1 
times normal in concentration. 




Fig. 



29 



84 Introduction to General Chemistry 

It is often convenient in practice to use solutions of J, -=-, T V, 
or some other fraction of normal; we call these half-normal 

— ), one-fifth normal (— ), and one-tenth or deci-normal ( - 

respectively. 

137. Acidimetry and Alkalimetry. — The analyses of acids 
and bases by means of normal solutions are called respectively 
acidimetry and alkalimetry. The act of running in a solution 
from a burette until the neutral or end-point is reached is called 
titration. The volume of solution used is called the titer. 
Instead of litmus we may use some other colored substance to 
indicate the end-point; such a substance is called an indicator. 
Other useful indicators are methyl orange, phenolphthalein, 
and Congo red. 

138. Problems. — 

N 

1. How many c.c. of — nitric acid are required to neutralize 

50 c.c. of normal potassium hydroxide? (107) 

N 

2. How many c.c. of — sodium hydroxide are required to 

N 
neutralize 20 c.c. of — sulfuric acid? (94) 

3. If 16 c.c. of a solution of sulfuric acid of unknown concen- 

N 
tration requires for its neutralization 36 c.c. of — potassium 

hydroxide, (a) what is the weight of sulfuric acid in the 16 c.c. 
taken? (b) what is the weight of sulfuric acid in 1 liter of this 
acid? (107) 

139. The Formation of Water. — We have, in earlier chapters, 
learned various ways in which water can be formed chemically. 
We may enumerate these by way of review. 

Water is formed — 

1. By the burning of hydrogen: 

2H 2 +0 2 -> 2 H 2 0. 

2. By the burning of a compound of hydrogen, for example, 

methane : 

CH 4 + 2 2 -> C0 2 + 2 H 2 0. (86) 



Water and Solutions 85 

3. By the oxidation of hydrogen or its compounds by means 
of combined oxygen, as, for example, when ammonia is passed 
over hot copper oxide: 

2NH3+3C11O -> 3 H 2 0+ 3 Cu+N 2 . (84) 

4. By the union of acids and bases, whereby a salt and water 
are always formed; for example: 

HCl+NaOH -> NaGl+H 2 0. (89) 

5. By the decomposition or dissociation of various unstable 
compounds, as, for example, sodium sulfate decahydrate into 
the anhydrous salt and water: 

Na 2 S0 4 • ioH 2 -> NaS0 4 + ioH 2 0. (96) 

140. The Chemical Reactions of Water. — We have also 
studied some of the important kinds of reactions in which water 
takes part. We may now summarize these as follows: 

1. Water unites with salts to form hydrates, thus: 

Na 2 S0 4 + ioH 2 -> Na 2 S0 4 . ioH 2 0. (96) 

2. Ammonia and water unite to form ammonium hydroxide: 

NH 3 +H 2 -> NH 4 OH. (91) 

4. Water acts upon some metals to give hydroxides and 
hydrogen. Thus, sodium and cold water react very easily, 
giving sodium hydroxide and hydrogen: 

2 Na+2H 2 0-> 2 NaOH+H 2 . (88) 

Magnesium does not act readily on cold water, but burns vigor- 
ously in steam giving the hydroxide and hydrogen : 

Mg+2H a O->.Mg(OH) a +H 2 . (28, 86) 



CHAPTER IX 
ACIDS, BASES, AND SALTS.— II 

141. New Acids, Bases, and Salts. — The present chapter will 
treat of three new acids, carbonic, H 2 C0 3 , acetic, C 2 H 4 2 , and 
phosphoric, H 3 P0 4 , and the bases derived from the elements 
magnesium, calcium, barium, zinc, iron, aluminum, copper, 
silver, lead, and mercury, together with the more important salts 
which these bases form with the three acids studied in the first 
chapter on acids, bases, and salts, as well as with the three acids 
above-mentioned. 

142. The Action of Water on Magnesium Oxide : Magnesium 
Hydroxide, Mg(0H) 2 . — All of the three bases studied in the first 
chapter on "Acids, Bases, and Salts" are readily soluble in water. 
We shall next consider one which dissolves in water only to a very 
slight extent. If we shake, with water, a little magnesium 
oxide (11, 80), obtained by burning magnesium, we find that the 
solution will turn red litmus blue, although but a small amount 
of the magnesium oxide has dissolved in the water, the larger 
part having remained undissolved. It has been found by careful 
experiment that magnesium oxide and water unite when brought 
together, giving a single new compound, the composition of 
which is represented by Mg0 2 H 2 , which we may also write 
Mg(OH) 2 , and call magnesium hydroxide. This is a white sub- 
stance, with which the student may already be familiar under the 
name of " milk of magnesia." It is extensively used in medicine. 
It is to be classified as a base, since, like sodium hydroxide, it 
colors litmus blue and neutralizes hydrocloric acid. The equa- 
tion for the action of water on the oxide is 

MgO+H 2 0->Mg(OH) 2 . 

143. The Action of Hydrochloric Acid on Magnesium Hy- 
droxide: Magnesium Chloride, MgCl 2 . — Magnesium hydroxide 
is but very slightly soluble in water. However, if we add hydro- 
chloric acid to the magnesium hydroxide formed from the mag- 

86 



Acids, Bases, and Salts — II 87 

nesium oxide and water until the solution just turns litmus red, 
we find that all of the solid dissolves, giving a clear, colorless 
solution which if left to evaporate in an open vessel will deposit 
colorless crystals. An investigation of this new substance shows 
that it is a compound of magnesium, chlorine, hydrogen, and 
oxygen, in the proportion indicated by the formula MgCl 2 *6H 2 0. 
This hydrate of magnesium chloride is formed as a result of the 
following two reactions: 

Mg(OH) 2 -f 2HCI -> MgCl 2 + 2 H 2 0, 
MgCl 2 +6H 2 -> MgCl 2 • 6H 2 0. 

144. Magnesium Sulfate, MgS0 4 . — If we now add diluted 
sulfuric acid to some magnesium hydroxide mixed with water, 
until all of the solid has dissolved and litmus shows the solution 
to be neutral, we may obtain from the solution by careful 
evaporation crystals of magnesium sulfate having the formula 
MgS0 4 '7H 2 0, a substance much used in medicine and known as 
Epsom salts. The reaction occurs according to the equation: 

H 2 S0 4 +Mg(OH) 2 -> MgS0 4 +2H 2 0, 

the MgS0 4 then combining with water from the solution, thus: 

MgS0 4 +7H 2 -> MgS0 4 -.7H 2 0, 

to form the hydrate. The latter, when heated, readily disso- 
ciates into MgS0 4 and 7H 2 0, a fact which may be expressed thus: 

MgS0 4 - 7H 2 -> MgS0 4 + 7H 2 0. 

This, as we see, is just the reverse of the preceding reaction. 
The reactions of hydrates in solution are of course the same as 
those of the anhydrous salts, since solutions of the two cannot 
be distinguished. In what follows, the discussion of the hydrates 
formed will be omitted except in a single important instance. 

145. Magnesium Nitrate, Mg(N0 3 ) 2 . — Magnesium hydroxide 
is readily neutralized by nitric acid, with the formation of 
magnesium nitrate, which forms white crystals very easily 
soluble in water: 

Mg(OH) 2 + 2HNO3 -> Mg(N0 3 ) 2 + 2H a O. 



88 Introduction to General Chemistry 

Magnesium oxide and dilute hydrochloric acid react to give 
magnesium chloride, which is the same compound as that formed 
from magnesium hydroxide and the same acid. The equation 
for the reaction is 

MgO+ 2HCI -> MgCl 2 +H 2 0. 

The corresponding reactions take place with sulfuric and with 
nitric acid, and are represented by the equations 

MgO+H 2 S0 4 -> MgS0 4 +H 2 0, 
MgO+2HN0 3 ^ Mg(N0 3 ) 2 +H 2 0. 

146. Monacid and Diacid Bases; Valence. — If we compare 
the formula of magnesium chloride, MgCl 2 , with that of sodium 
chloride, NaCI, or potassium chloride, KC1, we see that, in the 
first case, one symbol weight of the metal is combined with two 
symbol weights of chlorine, while in the other two cases one 
symbol weight of metal is combined with but one symbol weight 
of chlorine. In the cases of the neutralization of the hydroxides 
by hydrochloric acid we found that one formula weight of mag- 
nesium hydroxide required two formula weights of hydrochloric 
acid (143) ; while one formula weight of the hydroxide of either 
sodium or potassium required but one of hydrochloric acid (102, 
107). For this reason we call sodium and potassium hydroxides 
monacid bases, and magnesium hydroxide a diacid base. We 
also make use of the term valence in referring to facts like those 
just mentioned, saying that the valence of sodium or potassium 
is one, while that of magnesium is two, or that sodium and 
potassium are univalent, while magnesium is a bivalent element. 
Since hydrogen chloride has the formula HC1, we say that hydro- 
gen has a valence of one, and we also say that the valence of 
chlorine is one. 

147. Radicals and Their Valence. — We have already become 
acquainted with several ammonium compounds, as, for example, 
the chloride NH 4 C1 (91) and the nitrate NH 4 N0 3 (105). We 
call the combination NH 4 the ammonium radical; it has never 
been obtained as a separate substance, but is known only as a 
component of ammonium compounds. We know of many other 
such radicals, one of which is met with in sulfuric acid and sul- 



Acids, Bases, and Salts — II 89 

fates, where we have found that sulfur and oxygen are always 
present in the ratio represented by S0 4 . Here again we have a 
radical which is found in many salts, the sulfates, but is not 
known as a separate chemical substance. In nitric acid and 
the nitrates we have the radical N0 3 . A radical is composed of 
two or more elements united in a definite proportion; con- 
sequently the composition of a radical can always be represented 
by a formula. We may consider that the combination of nitro- 
gen and hydrogen, NH 4 , taken as a radical, has a valence of one, 
since the weight of nitrogen and hydrogen represented by NH 4 
taken once unites with one symbol weight of chlorine, giving 
NH4CI. Since sulfuric acid, H 2 S0 4 , forms such salts as Na 2 S0 4 , 
K 2 S0 4 and (NH 4 ) 2 S0 4 , we say that the sulfate radical, S0 4 , has 
a valence of two, a fact which is also shown by the existence of 
such salts as NaHS0 4 , etc. Now if magnesium has a valence 
of two and the sulfate radical has also the valence of two, we 
see in the fact that magnesium sulfate has the formula MgS0 4 , a 
broader meaning of the term valence. And so chemists often 
speak of the two valences of magnesium being satisfied by the 
two valences of the sulfate radical. The subject of valence will 
be considered again at the end of this chapter (183). 

148. Zinc and Its Salts. — Zinc is an element and a very 
important metal; it is known in commerce as spelter, and is used 
in enormous amounts in making galvanized iron, which is iron 
coated with metallic zinc, in making brass, whose other com- 
ponent is copper, and for many other purposes. Zinc will burn 
when strongly heated in the air or in oxygen, giving a white 
oxide, the reaction being represented by the equation 

2Zn+0 2 ->2ZnO. 

Zinc oxide is used extensively in making white paint. 

It will be recalled that magnesium burns, giving an oxide 
(11, 80), and that this oxide reacts with acids giving salts, thus: 

MgO+2HCl->MgCl 2 +H 2 0. (143, 145) 

Zinc oxide behaves like magnesium oxide when treated with 
hydrochloric acid, giving zinc chloride, thus: 
ZnO+2HCl-> ZnCL+H 2 0. 



90 Introduction to General Chemistry 

Zinc chloride is a salt which dissolves very readily in water, 
giving a clear, colorless solution. Zinc oxide gives zinc sulfate 
and zinc nitrate as follows: 

ZnO+H 2 S0 4 -> ZnS0 4 +H 2 0, 
ZnO+ 2HNO3 -> Zn(N0 3 ) 2 +H 2 0. 

These are white salts, also easily soluble in water. 

149. The Action of Hydrochloric Acid on Zinc. — If we pour 
some hydrochloric acid on zinc we observe a vigorous reaction; 
the zinc dissolves and a gas which proves to be hydrogen is given 
off. If, after the zinc has all dissolved, we evaporate the solution, 
we obtain a white solid which is found to be zinc chloride. The 
reaction is represented thus: 

Zn+ 2 HCl->ZnCl 2 +H 2 . 

Comparing this equation with that for the action of zinc oxide 
on hydrochloric acid, we see that in the latter case the hydrogen 
of the acid, instead of passing off as gas, unites with the oxygen 
of the zinc oxide, giving water. 

We might expect that metallic magnesium and hydrochloric 
acid would act thus: 

Mg+2HCl->MgCl 2 +H 2 , 

and it is easy to show by experiment that this is the case. With 
dilute sulfuric acid these metals behave as follows: 

Mg+H 2 S0 4 -> MgS0 4 +H 2 , 
Zn+H 2 S0 4 -> ZnS0 4 +H 2 . 

In making hydrogen in the laboratory we usually use zinc and 
hydrochloric acid. 

150. Marble and Other Compounds of the Element Cal- 
cium. — Let us now consider the chemical behavior of marble. 
If we place some lumps of marble in a hard glass tube and heat 
strongly, a gas is given off, while the lumps change but little in 
appearance. This gas causes limewater to turn milky; it is 
carbon dioxide (19). If the lumps left after heating the marble 
are moistened with water, they grow very hot, swell up, and 
crumble to a white powder. It is evident therefore that the 



Acids, Bases, and Salts — 77 91 

marble has been changed chemically by the heating. The solid 
left after the heating is the common substance, quicklime. The 
action of water upon quicklime is called slaking. If the slaked 
lime is shaken with a large amount of water, not much seems to 
dissolve; but if we filter the mixture, a clear, colorless solution 
is obtained. If some carbon dioxide gas is run into this clear 
solution it turns milky, because this solution is limewater (18), 
of which we have so often made use. If we test the limewater 
with litmus we find that it turns the latter blue, showing the 
limewater to be a solution of a base. This base reacts with acids 
to form salts. All of these products contain an element called 
calcium, whose symbol is Ca. Calcium is a brassy-looking metal, 
which will readily burn with a bright light if heated in air or 
oxygen, giving calcium oxide: 

2Ca+0 2 ->2CaO. 

Calcium oxide, CaO, is quicklime; but the latter is never 
made practically in this way, because metallic calcium is too 
expensive, and because the oxide is made very cheaply by heating 
marble or, more often, limestone, which is an impure form of the 
same compound as marble. By Seating marble, CaO and C0 2 
are formed, and nothing else. By finding the percentage of 
each we can easily calculate the formula for marble to be CaC0 3 , 
which is called calcium carbonate; the effect of the heating is, 
therefore, represented thus: 

CaC0 3 ->CaO+C0 2 . 

151. Calcium Hydroxide, Ca(OH) 2 . — As has been stated, 
when water acts on calcium oxide or quicklime, we get calcium 
hydroxide or slaked lime, a solution of which is called limewater: 

CaO+H 2 0->Ca(OH) 2 . 

This reaction is analogous to the action of water on magnesium 
oxide, which was studied earlier (142). The action of hydro- 
chloric acid on calcium hydroxide gives calcium chloride and 
water: 

Ca(OH) 2 +2HCl -> CaCl 2 +2H 2 0. 



92 Introduction to General Chemistry 

We are now in position to understand the cause of the milki- 
ness produced when carbon dioxide acts on limewater. The 
white solid formed is really calcium carbonate, CaC0 3 . The 
equation is 

Ca(OH) 2 +C0 2 -> CaC0 3 +H 2 0. 

152. Carbonic Acid, H 2 C0 3 . — If we pass carbon dioxide into 
water, a solution results which has faint acid properties. This 
solution is in fact the well-known plain soda served at soda 
fountains. The dissolved carbon dioxide and water partially 
combine to form an acid called carbonic acid : 

C0 2 +H 2 0->H 2 C0 3 . 

Therefore we may then consider that it is this acid which 
neutralizes the base calcium hydroxide, thus : 

Ca(OH) 2 +H 2 C0 3 -> CaC0 3 + 2 H 2 0. 

Calcium carbonate is a salt which is almost insoluble in water. 
In fact, salts exhibit all degrees of solubility in water. Some, 
like zinc chloride, dissolve in less than their own weight of water; 
others, like common salt, are much less soluble; while many, like 
calcium carbonate, are very nearly insoluble in water. 

153. Calcium Sulfate, CaS0 4 . — Calcium hydroxide and sul- 
furic acid form calcium sulfate and water: 

Ca(OH) 2 +H 2 S0 4 ^ CaS0 4 + 2 H 2 0. 

We find by experiment that the calcium sulfate so formed dis- 
solves very slightly in water, 100 c.c. of water dissolving but 
one-fourth of a gram of the salt. On the other hand, calcium 
chloride is very soluble in water. If we add to a solution con- 
taining, say, 5 or 10 per cent of calcium chloride, a sufficient 
amount of sulfuric acid, we observe that a large amount of a 
white powder forms in the mixed solutions and soon settles to 
the bottom of the vessel, leaving a clear, colorless liquid above. 
The white powder proves to be calcium sulfate, which^is formed 
thus: 

CaCl 2 +H 2 S0 4 -> CaS0 4 + 2 HCl. 



Acids, Bases, and Salts — II 93 

154. Precipitation. — We often encounter chemical reactions 
in which, as in the action between calcium chloride and sulfuric 
acid, a solid is formed upon bringing together two solutions. A 
solid so thrown down is called a precipitate ; and we speak of the 
precipitation of calcium sulfate. The formation of insoluble 
calcium carbonate by the action of carbon dioxide on limewater 
is another example of precipitation. 

155. Gypsum and Plaster of Paris. — Calcium sulfate occurs 
in nature as the mineral gypsum, CaS0 4 *2H 2 0, which, if the 
water of hydration is driven off by heat, is converted into the 
well-known plaster of Paris: 

CaS0 4 - 2H 2 -> CaS0 4 + 2H 2 0. 

When powdered plaster of Paris is mixed with enough water to 
form a paste, it sets in the course of an hour into a solid mass 
which retains the form of the vessel or mold which holds it. 
Plaster casts are made in this way. The setting is due to the 
formation of interlacing crystal filaments of the hydrate 
CaS0 4 * 2H 2 0, formed by a reversal of the action by which plaster 
of Paris is formed from gypsum: 

CaS0 4 +2H 2 -> CaS0 4 - 2H 2 0. 

156. Calcium Bicarbonate and Hard Water. — A very inter- 
esting and important reaction occurs when carbon dioxide is 
passed for a long time into a sufficiently dilute solution of calcium 
hydroxide (limewater). At first a milkiness appears, due to the 
formation of calcium carbonate : 

Ca(OH) 2 + C0 2 -> CaC0 3 +H 2 0. 

If we continue to pass in carbon dioxide, the precipitate slowly 
dissolves, giving finally a perfectly clear solution. If this solu- 
tion is now boiled, carbon dioxide gas is given off and a white 
precipitate is formed. These facts are explained in the following 
way. Carbonic acid, H 2 C0 3 , like sulfuric acid, is a dibasic 
acid (102) and can form acid salts as well as neutral salts. Just 
as sulfuric acid yields Na 2 S0 4 and NaHS0 4 , so carbonic acid gives 
Na 2 C0 3 and NaHC0 3 , sodium carbonate and sodium acid car- 
bonate, also known as bicarbonate (baking-soda). 



94 Introduction to General Chemistry 

The calcium salts corresponding to sodium carbonate and 
bicarbonate are CaC0 3 and Ca(HC0 3 ) 2 . The difference in the 
formulae of the sodium and calcium salts is due to the fact that 
the valence of calcium is two, while that of sodium is one. Now 
when carbon dioxide, in excess, acts on calcium carbonate, 
calcium acid carbonate, called also bicarbonate, is formed, and 
this being soluble in water the precipitate goes into solution: 

C0 2 +H 2 0->H 2 C0 3 
CaC0 3 +H 2 C0 3 -> Ca(HC0 3 ) 2 . 

When the clear solution so obtained is boiled the following reac- 
tion occurs: 

Ca(HC0 3 ) 2 -> CaC0 3 +H 2 0+C0 2 . 

These reactions take place extensively in nature. Natural 
waters, e.g., those of springs and rivers, contain dissolved carbon 
dioxide, and therefore carbonic acid. Such waters passing over 
limestone, impure CaC0 3 , dissolve it and take the Ca(HC0 3 ) 2 
into solution, forming so-called hard water. When boiled, as 
in a teakettle, it gives off carbon dioxide and deposits the calcium 
carbonate. 

157. Vinegar: Acetic Acid, C 2 H 4 2 . — Acetic acid is the 

principal ingredient, other than water, in vinegar, of which it 

constitutes about 4 per cent. The formula of acetic acid is 

C 2 H 4 2 . It neutralizes sodium hydroxide according to the 

equation 

C 2 H 4 2 +NaOH = NaC 2 H 3 2 +H 2 0. 

The salt NaC 2 H 3 2 , sodium acetate, is the only sodium salt 
which can be made from this acid. Therefore the acid radical 
of acetic acid and its salts is C 2 H 3 2 and we may write the 
formula of the acid HC 2 H 3 2 to indicate that only one of the four 
hydrogen atoms of a molecule is replaceable in salt formation. 

Pure acetic acid is a colorless liquid, miscible with water in 
all proportions. It is monobasic and forms with most bases salts 
called acetates. 

158. Bone Ash: Calcium Phosphate, Ca 3 (P0 4 ) 2 . — When 
bones are burned only the gelatinous matter and connective 
tissue are removed; the white material which is left is called 



Acids, Bases, and Salts — // 95 

bone ash and consists essentially of calcium phosphate, Ca 3 (P0 4 ) 2 . 
If powdered bone ash, which is practically insoluble in water, is 
stirred with somewhat diluted sulfuric acid, the following reac- 
tion occurs: 

Ca 3 (P0 4 ) 2 + 3 H 2 S0 4 -> 3 CaS0 4 +2H 3 P0 4 . 

The calcium sulfate formed is difficultly soluble in water and may 
be filtered out, giving a clear, colorless filtrate containing dis- 
solved phosphoric acid, H 3 P0 4 . 

159. Phosphoric Acid: a Tribasic Acid. — This acid is a white 
crystalline solid, which is very soluble in water, frequently com- 
ing on the market in the form of a very concentrated solution of 
syrupy consistency. Its dilute solution has a pleasant sour taste 
and turns litmus red. With suitable proportions of sodium 
hydroxide it yields the three.salts, Na 3 P0 4 , trisodium phosphate, 
Na 2 HP0 4 , disodium hydrogen phosphate, and NaH 2 P0 4 , 
sodium dihydrogen phosphate. The latter is a typical acid salt, 
having a sour taste and acid action on litmus. Phosphoric acid 
is therefore a tribasic acid. It forms with bases three series of 
salts, corresponding to those of sodium. To distinguish these 
classes of salts from one another they are called primary, second- 
ary, and tertiary, that with the smallest proportion of base being 
the primary and that in which all hydrogen is replaced being the 
tertiary. 

160. The Practical Importance of Calcium Phosphate. — 
Since calcium phosphate, Ca 3 (P0 4 ) 2 , constitutes the mineral 
matter of bones, it is, of course, a substance of very great im- 
portance. Phosphates in small amounts are also indispensable 
constituents of most plants, and it is from these, especially from 
the seeds, like wheat, oats, and corn, that men and animals get 
their needed supply. Plants, in turn, get their phosphates from 
the soil, and do not thrive on soil deficient in phosphates. Such 
infertile soil may be greatly improved by the use of fertilizers 
containing phosphates. For this purpose, bone ash is often 
employed; but since bone ash is almost insoluble in water, it is 
not directly available for plant use. In order to make it available 
it is treated with sufficient sulfuric acid to convert it into 



96 Introduction to General Chemistry 

Ca(H 2 P0 4 ) 2 , usually known as calcium superphosphate, which 
is soluble in water : 

Ca 3 (P0 4 ) 2 +2H 2 S0 4 ->Ca(H 2 P0 4 ) 2 +2CaS0 4 . 

Immense deposits of calcium phosphate occur in Florida and 
Tennessee, as phosphate rock. These deposits have doubtless 
been formed in past geological ages from the bones of marine 
animals. Phosphate rock, after treatment with sulfuric acid as 
in the case of bone ash, is used in enormous quantities as a 
fertilizer. 

161. Sodium Carbonate and Bicarbonate. — The carbonates 
of sodium which were referred to above (156) may be obtained 
by passing carbon dioxide into sodium hydroxide solution; we 
get in this way either the carbonate Na 2 C0 3 , or the bicarbonate 
NaHC0 3 , according to the proportion of carbon dioxide used. 
We may consider that the gas first unites with water to form 
carbonic acid, which then reacts with sodium hydroxide according 
to the two following equations: 

2 NaOH+H 2 C0 3 -> Na 2 C0 3 + 2 H 2 0, 
NaOH+H 2 C0 3 H> NaHC0 3 +H 2 0. 

These carbonates of sodium are manufactured in immense 
quantities, as they are very important substances. In practice 
they are not made according to the reactions given, but by more 
economical processes, which will be considered later. 

162. Potassium Carbonate and Bicarbonate. — Potassium also 
forms analogous carbonates, K 2 C0 3 and KHC0 3 ; the former, 
commonly known as potash, is contained in wood ashes, from 
which it may be dissolved by water. Upon boiling down the 
solution known popularly as lye, a residue of crude potassium 
carbonate, K 2 C0 3 , remains. This, when more strongly heated 
to burn out brown tarry matters, gives white potash, so called 
from the fact that the evaporation of the lye is carried out in an 
iron pot. This lye is extensively used in the preparation of a 
crude soft soap. A purer form of potash is used in manufacturing 
liquid soaps. Common hard soap is made from sodium car- 
bonate and fats of various kinds. 



Acids, Bases, and Salts — II 97 

163. The Action of Acids on Carbonates. — If some hydro- 
chloric acid is poured on a piece of marble (150), the liquid 
appears to boil, although the temperature does not rise notice- 
ably. It is easy to show that the apparent boiling, called 
effervescence, is due to the escape of carbon dioxide gas. The 
marble dissolves completely if sufficient acid is used, and the 
evaporated solution leaves a residue of calcium chloride (151). 
The reaction is as follows: 

CaC0 3 +2HCl-> CaCl 2 +H 2 0+C0 2 . 

Similar reactions take place between calcium carbonate and 
nitric and sulfuric acids: 

CaC0 3 +2HN0 3 ^ Ca(N0 3 ) 2 +H 2 0+C0 2 , 
CaC0 3 +H 2 S0 4 -> CaS0 4 +H 2 0-f-C0 2 . 

In fact, the carbonates of other elements all show this kind of a 
reaction with these acids; for example: 

NaHC0 3 +HCl -> NaCl+H 2 0+ C0 2 , 
K 2 C0 3 + 2 HN0 3 ^ 2KN0 3 -f-H 2 0+C0 2 . 

In general, carbonates are decomposed by acids. 

164. Barium Sulfate: a Test for Sulfates. — The element 
barium resembles calcium (150) very closely in its behavior. 
Let us consider just one of its reactions at present, leaving a 
study of the others until a later time. Barium sulfate, BaS0 4 , 
is a white solid which is as insoluble in water as glass; barium 
chloride, BaCl 2 , is about as soluble as common salt. If we pour 
some sulfuric acid into a clear, colorless solution of barium 
chloride, a white precipitate of barium sulfate forms at once : 

BaCl 2 +H 2 S0 4 -> BaS0 4 +2HCl. 

We should observe the similarity of this equation to that for the 
action of sulfuric acid on calcium chloride (153). 

If we add to a solution of barium chloride a solution of sodium 
sulfate, or of magnesium sulfate, or in fact of any sulfate whatso- 
ever, a precipitate of barium sulfate is formed. For example, 
with magnesium sulfate we have 

BaCl 2 +MgS0 4 -> BaS0 4 +MgCL. 



98 Introduction to General Chemistry 

By means of this reaction we can tell at once whether any solu- 
tion contains sulfuric acid or a sulfate: if no white precipitate 
is formed, sulfuric acid and sulfates are absent. We call this 
a test for sulfuric acid or sulfates. 

165. Copper and Its Compounds. — The important, familiar 
metal copper is an element. We have already learned that when 
heated in air or oxygen it unites with oxygen to form copper 
oxide (32, 33, 82), a black solid: 

2Cu+0 2 ->2CuO. 

Copper oxide reacts with the corresponding acids to form the 
chloride, nitrate, and sulfate, thus: 

CuO+2HCl-> CuCl 2 +H 2 
CUO+2HNO3 -> Cu(N0 3 ) 2 +H 2 
CuO+H 2 S0 4 -> CuS0 4 +H 2 0. 

These salts are all easily soluble in water, giving blue solutions. 
It will be* recalled that calcium oxide unites with water to 
form the hydroxide, thus: 

CaO+H 2 0->Ca(OH) 2 . 

On the other hand, if we bring copper oxide and water together, 
no union takes place. This might be taken to indicate that 
copper hydroxide, which we might expect to have the formula 
Cu(OH) 2 , cannot be formed or does not exist. This, however, 
is not the case; it is a well-known substance which is easily 
obtained in another way. If we add to a solution of copper sul- 
fate a solution of sodium hydroxide, a blue precipitate of copper 
hydroxide forms. This is a blue solid which is very nearly 
insoluble in water. Its formation takes place thus: 

CuS0 4 +2NaOH-> Cu(OH) 2 +Na 2 S0 4 . 

We may also get copper hydroxide by the interaction of solutions 
of copper chloride or nitrate with sodium hydroxide: 

CuCl 2 + 2 NaOH -> Cu(OH) 2 + 2 NaCl. 

If the copper hydroxide formed in the last reaction is heated 
by boiling the mixture, the blue precipitate turns black. This 



Acids, Bases, and Salts — 77 99 

change in color is due to a change of part of the hydroxide into 
the oxide and water: 

Cu(OH) 2 ->CuO+H 2 0. 

166. The Preparation of Difficultly Soluble Hydroxides. — 

Many hydroxides of elements are nearly insoluble in water. In 
such cases, the hydroxides are formed from a solution of a salt 
of the element by adding sodium hydroxide, or potassium 
hydroxide, or in many cases ammonium hydroxide, as illustrated 
by the following equations : 

CaCl 2 + 2 KOH-> Ca(OH) 2 +2KCl, 
MgS0 4 + 2 NaOH -> Mg(OH) 2 +Na 2 S0 4 . 

Such difficultly soluble hydroxides separate as precipitates, which 
may be filtered out. 

167. Lead and Its Compounds. — The well-known metal lead 
is an element, which is used extensively in metallic form and also 
in the form of compounds. Lead unites with oxygen directly 
when heated in air or oxygen, giving, under suitable conditions, 
the pale yellow oxide PbO, known as litharge. This oxide, like 
those of magnesium, zinc, and copper, reacts with acids to form 
salts. Thus with nitric acid we get lead nitrate : 

PbO+ 2HNO3 ~> Pb(N0 3 ) 2 +H 2 0. 

This salt forms large, white crystals which dissolve readily in 
water to form a colorless solution. If hydrochloric acid is added 
to a solution of lead nitrate, a white precipitate of lead chloride 
is obtained: 

Pb(N0 3 ) 2 + 2HCI -> PbCl 2 + 2HNO3. 

Lead chloride is only slightly soluble in cold water, but is much 
more soluble in hot water, from which, upon cooling, it separates 
again in white needle-shaped crystals. Lead chloride is also 
obtained from litharge and hydrochloric acid: 

PbO+ 2HCI -» PbCl 2 +H 2 0. 



ioo Introduction to General Chemistry 

Upon adding dilute sulfuric acid to a solution of lead nitrate, 
a white precipitate of lead sulfate forms : 

Pb(N0 3 ) 2 +H 2 S0 4 -> PbS0 4 + 2HNO3. 

Lead sulfate is nearly insoluble in hot or cold water. 

Metallic lead is acted upon very slowly by hydrochloric acid. 
With the cold dilute acid no appreciable action takes place ; with 
boiling, concentrated acid, a very slow reaction occurs, thus: 

Pb+ 2 HCl^PbCl 2 +H 2 . 

By methods to be considered later, it is possible to prepare 
an oxide of lead containing double the proportion of oxygen 
present in PbO, namely Pb0 2 , or lead dioxide. This oxide does 
not react with dilute nitric acid. When it is heated with hydro- 
chloric acid it gives lead chloride and chlorine : 

Pb0 2 +4HC1 -> PbCl 2 + 2 H 2 0+ Cl 2 . 

If we compare this reaction with the following, 

PbO+ 2HCI -> PbCl 2 +H 2 0, 

we see that the excess of oxygen in Pb0 2 above that in PbO oxidizes 
the hydrochloric acid, forming water and setting free chlorine. 

Lead acetate, Pb(C 2 H 3 2 ) 2 *3H 2 0, is formed by dissolving 
litharge, PbO, in acetic acid. It forms colorless prismatic crys- 
tals, which are readily soluble in water. It is a poisonous salt 
and is called sugar of lead on account of its sweetish taste. 

168. Silver and Its Compounds. — Silver is so familiar a 
metal that we need not describe its properties. It is an element 
which is most extensively used in the metallic form, but which 
forms several compounds of great practical importance. The 
metal is not readily acted upon by dilute acids, with the excep- 
tion of nitric acid, with which it undergoes a complex reaction 
represented by the equation 

3 Ag+4HN0 3 -> 3 AgN0 3 +NO+2H 2 0. 

We need not consider this reaction critically at this time, 
although it is well worth careful study; but note that silver 



Acids, Bases, and Salts — II 101 

nitrate is an easily soluble salt, forming a colorless solution. The 
solid salt forms large white crystals. 

169. Silver Chloride, AgCl. — The addition of hydrochloric 
acid to a solution of silver nitrate produces at once a heavy white 
precipitate of silver chloride, which is almost insoluble in water: 

AgN0 3 +HCl-> AgCl+HN0 3 . 

By adding an excess of hydrochloric acid practically all of the 
silver in a solution is precipitated. The precipitate does not 
dissolve appreciably in any of the common acids. It is, however, 
very easily soluble in ammonia solution, from which it is again 
thrown down if the solution is acidified with nitric or hydrochloric 
acid. If any solution of unknown nature gives with hydro- 
chloric acid a white precipitate which is insoluble in an excess of 
the acid, but easily soluble in ammonia, from which solution it is 
thrown down by acidifying the solution with hydrochloric acid, 
it is safe to conclude that the original solution contained a salt 
of silver. This series of reactions constitutes a test for silver in 
the form of a dissolved salt. 

170. Silver Sulfate, Ag 2 S0 4 . — This salt is formed as a white 
crystalline precipitate when sulfuric acid is added to a con- 
centrated solution of silver nitrate. It is not very soluble, 1 g. 
requiring about 200 c.c. of cold water for its solution. The same 
salt is also formed by the action of hot, concentrated sulfuric 
acid on metallic silver. 

171. Silver Phosphate, Ag 3 P0 4 . — This salt is formed as a 
yellow precipitate when sodium phosphate, Na 3 P0 4 , or some 
other soluble phosphate is added to a solution of silver nitrate: 

3 AgN0 3 +Na 3 P0 4 -> Ag 3 P0 4 + 3 NaN0 3 . 

This yellow precipitate is readily soluble in dilute nitric acid, 
forming a colorless solution. It also dissolves easily in aqueous 
ammonia, giving a colorless solution, from which it is again 
thrown down when the solution is exactly neutralized with nitric 
acid. 

Silver may be distinguished from lead most easily by reason 
of the solubility of lead chloride in hot water, in which silver 
chloride is insoluble. 



102 Introduction to General Chemistry 

172. Silver Oxide, Ag 2 0. — The addition of sodium hydroxide 

to a solution of silver nitrate gives a black precipitate of silver 

oxide, Ag 2 0. We might expect silver hydroxide, AgOH, to be 

formed thus: 

AgN0 3 +NaOH -> AgOH+NaN0 3 . 

Possibly this is what first happens, but, if so, the hydroxide 
formed changes at once into the oxide, 

2AgOH->Ag 2 0+H 2 0. 

It will be recalled that copper hydroxide is decomposed at the 
temperature of boiling water into the oxide and water (165). 
In the case of silver hydroxide the change takes place at room 
temperature. 

173. Iron and Its Compounds. — The element iron is the most 
important of all metals. It unites directly with oxygen at a red 
heat, forming the oxide Fe 3 4 (81). It can also form two other 
oxides, FeO and Fe 2 3 . The oxide Fe 3 4 is magnetic and is 
called magnetic iron oxide; FeO is called ferrous oxide (from 
ferrum, iron), while Fe 2 3 is called ferric oxide. Ferrous oxide 
gives, with the corresponding acids, ferrous chloride, FeCl 2 , and 
ferrous sulfate, FeS0 4 . These salts are also formed from iron 
by the following reactions: 

Fe+2HCl->FeCl 2 +H 2 , 
Fe+H 2 S0 4 ->FeS0 4 -fH 2 . 

In all ferrous compounds the valence of iron is two. 

The action of hydrochloric acid on ferric oxide takes place 

thus: 

Fe 2 3 +6HCl-> 2FeCl 3 + 3 H 2 0. 

The salt FeCl 3 is called ferric chloride. It is a dark-yellow 
substance which dissolves easily in water to form a yellow solu- 
tion. On the other hand, ferrous chloride, FeCl 2 , is pale green 
and forms a pale-green solution. We cannot get FeCl 3 from 
iron and hydrochloric acid, but we do get the salt by the action 
of chlorine on ferrous chloride, 

2FeCl 2 -j-Cl 2 ->2FeCl 3 , 



Acids, Bases, and Salts — II 103 

or on iron, 

2Fe+3Cl 2 ->2FeCl 3 . 

In ferric chloride the valence of the iron is three. We also know 
ferric nitrate, Fe(N0 3 ) 3 , and ferric sulfate, Fe 2 (S0 4 ) 3 . There are, 
therefore, two series of iron salts — the ferrous, in which the valence 
of iron is two, and the ferric, in which the valence is three. 

We can obtain the two hydroxides of iron, both of which are 
nearly insoluble in water, by the action of sodium hydroxide on 
solutions of ferrous and ferric salts : 

FeCl 2 + 2 NaOH -> Fe(OH) 2 + 2 NaCl, 
FeCl 3 + 3 NaOH-> Fe(OH) 3 + 3 NaCl. 

Ferrous hydroxide is white if pure, but is usually obtained as a 
dirty-green precipitate; this is due to partial oxidation by the 
action of oxygen of the air, with which it readily unites. Ferric 
hydroxide is a brown precipitate. These hydroxides unite with 
acids to form salts: 

Fe(OH) 2 + 2 HCl -> FeCl 2 +2H 2 0, 
Fe(OH) 3 + 3 HCl-> FeCl 3 + 3 H 2 0. 

174. Aluminum and Its Compounds. — The common metal 
aluminum is an element. As is well known, the metal is not 
acted upon by air or water. It reacts easily with dilute hydro- 
chloric acid, giving aluminum chloride, A1C1 3 , and hydrogen: 

2Al+6HCl-> 2 A1C1 3 + 3 H 2 . 

Upon evaporation, the solution deposits white crystals of the 
compound A1C1 3 *6H 2 0. It is not possible to obtain the anhy- 
drous salt, A1C1 3 by heating these crystals, for the purpose of 
driving off water, since they decompose thus: 

2 A1C1 3 -6H 2 -> Al 2 3 +6HCl-f 3 H 2 0. 

The anhydrous chloride is formed by the action of dry chlorine 

gas on aluminum: 

2 A1+ 3 C1 2 ->2A1C1 3 . 



104 Introduction to General Chemistry 

Aluminum chloride is easily soluble in water, forming a colorless 
solution. This solution gives with ammonia a white precipitate 
of aluminum hydroxide, Al(OH) 3 , which is insoluble in water: 

A1C1 3 +3NH 4 0H -> A1(0H) 3 + 3 NH 4 C1. 

When heated, the hydroxide gives the oxide and water: 

2A1(0H) 3 ->A1 2 3 + 3 H 2 0. 

Rubies and sapphires are natural forms of aluminum oxide. 
Emery, which is a valuable abrassive, is an impure form of the 
same substance. 

175. Various Aluminum Salts. — The hydroxide is a base 
which reacts with acids to give the corresponding salts, thus: 

Al(OH) 3 + 3 HCl-> A1C1 3 + 3 H 2 0, 
Al(OH) 3 + 3 HN0 3 -> A1(N0 3 ) 3 + 3 H 2 0, 
2 A1(0H) 3 + 3 H 2 S0 4 -> A1 2 (S0 4 ) 3 +6H 2 0. 

The nitrate and sulfate are easily soluble in water, giving color- 
less solutions. The well-known substance alum is potassium, 
aluminum sulfate, KAl(S0 4 ) 2 »i2H 2 0. It is obtained in large, 
colorless crystals when a solution made from potassium sulfate 
and aluminum sulfate is allowed to evaporate. The correspond- 
ing sodium and ammonium salts are well known, and have the 
formulae NaAl(S0 4 ) 2 -i2H 2 and NH 4 Al(S0 4 ) 2 -i2H 2 0, respec- 
tively. All such compounds are known as double salts ; chemists 
are familiar with a great variety of these. Other examples of 
well-known double salts are ammonium ferrous sulfate, (NH 4 ) 2 
Fe(S0 4 ) 2 -6H 2 0, and potassium cupric chloride, K 2 CuCl 4 *2H 2 0. 

176. Acid Reaction of Aluminum Salts. — Solutions of the 
chloride, nitrate, and sulfate of aluminum, and also of alum, are 
not neutral, as we might expect, but are distinctly acid in reaction. 
They also have a sour taste. On the other hand, we find that 
moist aluminum hydroxide, if it has been carefully washed free 
from the ammonia used in precipitating it, has no action on either 
blue or red litmus. It is also tasteless. Nevertheless, we call 
the hydroxide a base, because it unites with acids to form salts. 
We say, however, that it is a weak base ; and we find in general 



Acids, Bases, and Salts — // 105 

that weak bases, of which many are known, give salts whose 
solutions are acid in reaction. This is an important matter which 
will have to be studied carefully later. 

177. Acid Properties of Aluminum Hydroxide. — If we add 

sodium hydroxide to a solution of an aluminum salt, a white 

precipitate of aluminum hydroxide is first formed, just as with 

ammonia : 

AlCl 3 + 3 NaOH-> Al(OH) 3 + 3 NaCl. 

However, upon adding an excess of sodium hydroxide, we find 
that the precipitate goes into solution. If pure aluminum hy- 
droxide is dissolved in a solution of sodium hydroxide and the 
resulting solution evaporated, crystals of sodium aluminate, 
NaA10 2 , are obtained. This substance is easily soluble in water 
and is, in reality, a salt. It thus appears that aluminum hy- 
droxide acts as an acid in this case, and we might write the equa- 
tion for the action of. sodium hydroxide upon it thus : 

HA10 2 -H 2 0+NaOH->NaA10 2 +2H 2 0. 

We find that the solution of sodium aluminate is strongly alkaline 
toward litmus, and say, therefore, that, although aluminum 
hydroxide has some acid properties, it is a very weak acid. 

Thus we see that a substance may be both a base and an acid. 
Such a substance is said to be amphoteric. Several metallic 
hydroxides are amphoteric. Thus zinc hydroxide, Zn(OH) 2 , 
forms with hydrochloric acid, ZnCl 2 , and with sodium hydroxide, 
Na 2 Zn0 2 , sodium zincate. It is of interest to note that aluminum 
hydroxide does not react with carbonic acid, and in fact no car- 
bonate of aluminum has ever been made. Now, carbonic acid 
is a very weak acid, and aluminum hydroxide is a very weak base. 
In general, we find that very weak bases do not form salts with 
very weak acids. 

Compounds of aluminum are very abundant in the earth. 
Common clay and numerous kinds of common rocks are com- 
pounds of aluminum. 

178. Mercury and Its Compounds. — We have already learned 
something of the chemical behavior of mercury and mercuric 
oxide, HgO (13, 14, 86). The oxide, which is insoluble in water, 



106 Introduction to General Chemistry 

dissolves in dilute hydrochloric acid, giving mercuric chloride, 
HgCl 2 , and in nitric acid, giving mercuric nitrate, Hg(N0 3 ) 2 : 

HgO+ 2HCI -> HgCl 2 +H 2 0. 
HgO+ 2HNO3 ~> Hg(N0 3 ) 2 +H 2 0. 

These salts form white crystals which are soluble in water. The 
soluble salts of mercury are all extremely poisonous when taken 
internally. Mercuric chloride is familiarly known as bichloride 
of mercury or corrosive sublimate, and is extensively used as a 
powerful germicide and antiseptic. 

179. The Formation of Mercuric Salts. — The nitrate can be 
made by the action of warm, concentrated nitric acid upon 
metallic mercury: 

3 Hg+8HN0 3 -> 3 Hg(N0 3 ) 2 +2NO+4H 2 0. 

Hydrochloric acid does not act appreciably upon mercury; but 
the chloride can be obtained by the action of chlorine on the 
metal : 

Hg+Cl 2 ->HgCl 2 . 

It is also made by heating a mixture of mercuric sulfate and 
common salt: 

HgS0 4 +2NaCl -> HgCl 2 +Na 2 S0 4 . 

The mercuric chloride formed is readily volatile and is separated 
by sublimation; hence the old name " corrosive sublimate.' 
The process of vaporization of a solid and the condensation of its 
vapor directly to the crystalline form is called sublimation. 
The sulfate is made by strongly heating mercury with concen- 
trated sulfuric acid: 

Hg+ 2 H 2 S0 4 -> HgS0 4 + S0 2 + 2H 2 0. 

180. Mercurous Salts. — The action of cold, dilute nitric acid 
on an excess of mercury gives rise to a solution of a salt having the 
formula HgN0 3 and called mercurous nitrate. A solution of 
this salt gives with hydrochloric acid a white precipitate of 
mercurous chloride, HgCl, and with dilute sulfuric acid also a 
white precipitate of mercurous sulfate, Hg 2 S0 4 ; both of these 



Acids, Bases, and Salts — II 107 

precipitates are practically insoluble in water. Thus mercury, 
like iron, forms two series of salts: the mercurous, in which the 
element has a valence of one, or is univalent, and the mercuric, 
in which it has a valence of two, or is bivalent. 

181. The Two Oxides of Mercury. — A solution of mercuric 
nitrate gives with a solution of sodium hydroxide a yellow 
precipitate of mercuric oxide, HgO: 

Hg(N0 3 ) 2 + 2 NaOH -> HgO+ 2NaN0 3 +H 2 0. 

The hydroxide of mercury, like that of silver, cannot be obtained; 
we might say that it is so unstable that it changes into the oxide 
and water as soon as it is formed; in this respect it resembles 
the corresponding compound of silver. The yellow oxide, 
formed in this way, seems to differ from the red oxide, obtained 
by heating mercury in the air Or in oxygen, only in being made 
up of very much smaller particles. 

A solution of mercurous nitrate gives with sodium hydroxide 
a nearly black precipitate of mercurous oxide, Hg 2 0. 

182. Calomel. — Mercurous chloride, HgCl, which is com- 
monly called calomel, is extensively used in medicine. It is a 
remarkable but well-known fact that the usual medicinal dose 
of calomel contains many times as much mercury as does a fatal 
dose of mercuric chloride. This great difference in physio- 
logical effect is in part due to the fact that while mercuric 
chloride is easily soluble in water, mercurous chloride is nearly 
insoluble. 

183. The Valencies of Radicals. — Now that we have studied 
a considerable additional number of acids, bases, and salts, 
we may again revert to the study of valence (146, 147) since it 
furnishes a key to the easy mastery of formulae, an undertak- 
ing which is as necessary to the study of chemistry as learning to 
spell is in the mastery of a language. To write the formula of a 
chloride of a metal it is only necessary to group with the s)inbol 
of the latter as many chlorine symbols as the metal in question 
has units of valence; thus the formulae of barium chloride and 
of aluminum chloride are BaCL and A1C1 3 respectively. If we 
wish to write the formulae of the nitrates we group the nitrate 



io8 



Introduction to General Chemistry 



radical with the metal symbol in question according to the same 
rule, thus Ba(N0 3 ) 2 and A1(N0 3 ) 3 . To write the formulae of 
sulfates, we must again group the symbols of the radical and 
the metal so that the total valence of each satisfies that of the 
other. Since the sulfate radical is bivalent, we may have to 
use more than one symbol weight of either the sulfate or the 
metal. Thus the formula of sodium sulfate is Na 2 S0 4 ; that of 
barium sulfate is BaS0 4 ; while that of aluminum sulfate is 
A1 2 (S0 4 ) 3 . Since P0 4 is trivalent, we know at once that the 
formula of sodium phosphate must be Na 3 P0 4 , that of barium 
phosphate must be Ba 3 (P0 4 ) 2 , and that of aluminum A1P0 4 . 
These examples are sufficient to show how a knowledge of 
valence simplifies the writing of formulae. In Table IX the 

TABLE IX 





H 


CI 
N0 3 

OH 
C 2 H 3 2 


HCl 
NaCl 

HC 2 H 3 2 
NH 4 C1 

AgCl 
HgCl 


HOH (H 2 0) 

HXO3 

NaOH 


I 


Na 

K 


Univalent 


NH 4 

Ag 


NH 4 OH 

Ag 2 




Hg 




Na 2 S0 4 










II 


Mg 

Ca 

Ba 

Zn 


S0 4 
C0 3 



MgCl 2 

CaCl 2 
BaCl 2 
ZnCl 2 


H 2 S0 4 
H 2 C0 3 
CaC0 3 
BaSO. 


Bivalent 


Fe 




FeCl 2 C11SO4 




Cu 




CuCl 2 J Ca(OH) 2 




Pb 




PbCl 2 MgO 

HgCl 2 ! HgO 




Hg 










III 

Trivalent 


Al 

Fe 


P0 4 

N 


AICI3 
FeCl 3 


H ? P0 4 
A1P0 4 
NH 3 


IV 

Quadrivalent 




C 


CC1 4 
C0 2 


CH 4 









various elements and radicals studied are classified with respect 
to their valencies, which vary from one to four. In the column 
headed by hydrogen we have the metals together with the 
ammonium radical. In the next column, headed by chlorine, 
are the elements and radicals that unite as a rule with 
those of the first column. The last two columns contain the 
formulae of some, typical compounds. 



CHAPTER X 

THE KINETIC THEORY OF MATTER AND THE MOLECULAR 
HYPOTHESIS 

184. An Old Greek Hypothesis. — The question whether a 
portion of a given substance, say a drop of water, could be sub- 
divided to an unlimited extent, and whether the smallest particle 
so produced would still differ in no way except in size from the 
original, was one which was much debated by the Greek philos- 
ophers centuries before chemistry became a science. Anaxagoras 
(b.c. 500), who held that there was no limit to the divisibility 
of matter, was opposed by Democritus (b.c. 470), who taught 
that in the imagined process of continued subdivision minute 
particles would finally be encountered which could not be cut 
in two without destroying or completely changing the nature 
of the substance; these particles were called atoms (aronos, a 
body which cannot be cut in two) . This idea may be illustrated 
by the following analogy. If we take a bushel of wheat, we may 
divide it into pecks, quarts, gills, etc., and yet each measure of 
the material will be a quantity of wheat; we may go still farther, 
but we will ultimately reach the single grains, which are still 
grains of wheat; but if we cut these grains in two the resulting 
parts may no longer be called wheat, since they would no longer 
possess the most remarkable property of wheat, which is that of 
growing if planted. At present we use the term molecule (from 
Latin molecula, the diminutive of moles , a mass) to mean essen- 
tially the same as the term "atom," as used by the Greeks, and 
speak therefore of the molecular theory of matter and the molec- 
ular hypothesis. The present chapter gives an account of 
this hypothesis and aims to show how it furnishes us an explana- 
tion of many important facts. 

185. The Molecules of Water. — According to the molecular 
hypothesis, a drop of water can be subdivided, and still remain 
water, only until the single molecules are reached, and no further. 
The splitting up of the molecules of water might separate the 

109 



no Introduction to General Chemistry 

smallest particles of oxygen from those of hydrogen; but the 
result would be the decomposition of the water into its elementary 
constituents. The particles of the elements oxygen and hydro- 
gen of which the molecules of water are made up are now called 
atoms. It is supposed that all of the molecules of water are 
alike in every respect, each being made up of one atom of oxygen 
and two atoms of hydrogen. The reason for this last conclusion 
is discussed in the next chapter. In general, each molecule of a 
given pure substance is just like every other molecule of that sub- 
stance. The nature of a substance is determined by the nature 
of its molecules; and this is, in turn, determined by the number 
and kind of atoms composing the molecules. We imagine mole- 
cules to be very small, since they cannot be seen with the aid 
of a microscope of the highest power. A cubic centimeter of air 
may contain an almost inconceivable number of these tiny 
particles. 

1 86. The Molecular Hypothesis Applied to Gases. — Let us 
first consider the known facts concerning gases and try to see how 
these facts can be connected with the supposition that gases are 
made up of small particles, the molecules. In the first place, 
we know that a confined gas tends to expand and exerts a pressure 
on the walls of the vessel which holds it. If we increase the 
pressure, the volume of the gas is diminished, and by applying 
a great pressure the decrease in volume may be made very great. 
We may explain this in either of two ways : first, that the mole- 
cules, like rubber balloons, are themselves compressible; or, 
second, that the molecules, which may not be appreciably com- 
pressible, are at considerable distances from one another, but 
are brought closer together when the gas as a whole is compressed. 
Let us follow up this second idea and see to what it leads. Two 
important questions now present themselves: (i) Why are the 
molecules not in contact — that is, why should they be at con- 
siderable distances from one another? (2) Why does a gas exert 
a pressure in all directions on the walls of the vessel which con- 
tains it? 

187. Are Molecules at Rest or in Motion?— Would it make a 
difference in the state of affairs whether the molecules were at 



Molecular Hypothesis in 

rest or in motion? Suppose they are in rapid motion: what 
would follow? Let us recall Newton's first law of motion: 
A body at rest remains at rest, and a body in motion continues to 
move with constant velocity in a straight line, unless acted upon 
by some external unbalanced force. Now, if molecules are in 
motion, and if they behave in the same manner as other bodies, 
and if, further, they are elastic — that is, if they, like rubber or 
ivory balls, can rebound from one another or from the walls of 
the containing vessel, then they will tend to continue in motion. 
Of course, in such a case as the one imagined, the molecules of the 
gas would very frequently strike one another and also the walls of 
the containing vessel; but, being elastic, they would rebound and 
continue in motion in a new direction; and although the velocity 
of an individual molecule might be increased or decreased as the 
result of a collision with another molecule, on the whole the 
average velocity of all the molecules would be constant. 

The various elaborations of the ideas here presented 
constitute the Kinetic Theory of Matter. This hypothesis is 
the most important corollary of the molecular hypothesis. It 
has proved enormously fruitful in explaining very diverse 
phenomena and in suggesting new lines of investigation. 

188. The Cause of Gas Pressure; Boyle's Law. — The strik- 
ing of a gas molecule against the wall of the vessel would deliver a 
little blow; and if millions of molecules struck each square centi- 
meter every second the effect would be to tend to push back the 
surface. But what is this but the exertion of pressure ? If the 
molecules strike often enough, regularly enough, and close enough 
together, this pressure would seem constant and uniform. Now, 
suppose the gas to be compressed until it occupies half its original 
volume. In each cubic centimeter there would now be double 
the original number of molecules, and on each square centimeter 
of the wall of the container twice as many molecules would 
strike per second as before; so that, as the mass of each mole- 
cule and also its velocity have remained unchanged, we should 
expect just double the pressure per square centimeter; and, in 
fact, this is just what we find by experiment. Thus we see that 
by imagining a gas to be made of numerous small, rapidly moving, 



ii2 Introduction to General Chemistry 

elastic particles, the molecules, we get an explanation of gas 
pressure and of Boyle's law. We also see how it would be pos- 
sible for the molecules to be at considerable distances (com- 
pared with their own diameters) from one another without 
tending to fall together into a mass in which the molecules would 
all be permanently in contact. 

189. The Effect of Temperature on Molecular Velocity. — We 
may next consider why it is that the molecules are in motion and 
whether the average velocity of the molecules of a given gas 
can ever be changed. We know, of course, that, for a constant 
volume, the pressure exerted by a gas increases with rise of 
temperature, so that if we are to explain gas pressure as due to the' 
momenta (mass X velocity) of the molecules which strike the walls 
of the container, we must suppose that an increase of tempera- 
ture increases either the mass of. a molecule or its velocity, or 
both. Naturally, the number of collisions with the wall could 
not be increased unless the velocity increased. Now, it would 
seem more reasonable to think of rise of temperature as causing 
an increase in velocity than to think of it as causing an increase 
in mass; so we have only to imagine that a rise in temperature 
causes the average velocity of the molecules to increase in order to 
get a simple and satisfying explanation of the effect of tempera- 
ture on the pressure of a gas. 

On the other hand, a decrease in temperature is accom- 
panied by a decrease in pressure; and, indeed, the pressure is at 
all times proportional to the absolute temperature. This would 
imply that the pressure would be zero at the absolute zero of 
temperature. But zero pressure could only result if the mole- 
cules were completely at rest. We may suppose, therefore, that 
at absolute zero there is no molecular motion. A body when hot 
differs from the same body when cold only by reason of the more 
rapid motion of its molecules. In short, according to this way 
of looking at the matter, heat is merely the outward manifestation 
of molecular motion. 

190. The Mixing of Gases. — We have already learned 
(122-124) that, as a rule, liquids and solids do not form perfect 
mixtures (solutions) in all proportions; that is to say, a solid or a 



Molecular Hypothesis 113 

liquid will dissolve in a second liquid only to a limited extent. 
Not so with gases : every gas will form with any proportion of 
any other gas a perfectly uniform mixture. Air, for example, is 
a perfectly homogeneous mixture of several gases. Of course, if 
some of the gases which are brought together react chemically, 
new liquid or solid compounds might be formed which would 
separate from the gaseous mixture. But for gases that do 
not react chemically we find that all gases mix perfectly in all 
proportions. 

191. The Diffusion of Gases. — If we bring two gases into the 
same vessel without attempting to mix them we find, after a time, 
that a perfectly uniform mixture is present in the vessel. We also 
know that if a gas like ammonia is liberated at one place in a 
closed room its odor is soon perceptible everywhere in the room. 
The process of the spontaneous mixing of gases is called diffusion, 
and we say that the ammonia has diffused through the air of the 
whole room. It is now easy to understand how this diffusion 
takes place. The molecules of ammonia are moving in all 
directions with high velocities; the same is true also of the mole- 
cules of the air; their complete and uniform intermingling is 
therefore inevitable. 

192. The Law of Partial Gas Pressures. — It is easily found by 
experiment that, if portions of various gases are brought together 
in the same vessel, the total pressure exerted by the gas mixture is 
the sum of the pressures that would be exerted, at the same tempera- 
ture, by the same portions of these gases if each occupied the space 
alone. This is known as Dalton's Law of Partial Pressures. 
It is not difficult to explain this law, since, in a mixture of gases, 
the molecules of a given sort will strike a given area of the Avail 
just as often in the presence of other unlike molecules as in their 
absence. Each kind of molecule will therefore produce the same 
partial pressure as if the others were absent. 

193. Avogadro's Hypothesis. — The student must already 
have been impressed by the fact that all gases show great similarity 
hi physical behavior. They all conform to the laws of Boyle and 
Charles. This fact, together with others which we shall consider 
later, led Avogadro, then professor of physics in Turin, Italy, to 



ii4 Introduction to General Chemistry 

suggest in 1811 that it is probable that all gases contain the same 
number of molecules per cubic centimeter. Although this sugges- 
tion received some support for the first twenty years after its 
proposal, it was then nearly forgotten until about i860, since 
which time its importance and probability have been impressed 
more and more deeply on the minds of physicists and chemists, 
so that during the last fifty years it has become one of the most 
fundamental principles of chemistry. It may be stated concisely 
thus : Equal volumes of every gas or vapor at the same temperature 
and pressure contain the same number of molecules. 

Further reasons for accepting Avogadro's hypothesis will 
be given in the next chapter. Indeed, the evidence from so 
many independent sources for the truth of this view is now so 
convincing that the hypothesis is looked upon by many as a 
statement of fact, and in consequence is referred to as Avogadro's 
Law. 

194. Gas Statistics. — Within the last few years methods have 
been found by means of which the number of molecules in 1 ex. 
of a gas has been found with a high degree of probability and 
accuracy. At o° and 76 cm., 1 c.c. of any gas has been found to 
contain 2 . 7 X 10 19 molecules, with a probable error of less than 1 
per cent. The number of molecules in 22.4 liters of a gas under 
standard conditions is therefore 2 2,4ooX2.7Xio I9 = 6.o6Xio 23 . 
The number of molecules in 1 c.c, twenty-seven millions of mil- 
lions of millions, is so immense that it is difficult for the mind to 
get any tangible conception of its magnitude. However, if we 
think of the molecules in 1 c.c. as at rest for the moment, and 
uniformly distributed in rows and layers, we should then have 
in each row of 1 cm. length 1^27Xio i8 = 3Xio 6 , or three million 
molecules, a number which, although large, is at least compre- 
hensible. Then in each layer there would be three million of 
these rows, and, in the whole cubic centimeter, three million such 
layers. 

195. A Cubic Mile of Sand. — Another mental picture of the 
case may be got if we imagine 1 c.c. of gas to have been expanded 
until it occupied a cubic mile. Then each row of molecules 
would be a mile long and would contain three million molecules, 



Molecular Hypothesis 115 

spaced about y V of an inch apart. Now, a grain of fine sand is 
about T V °f an mcn m diameter; and three million such grains 
placed side by side would extend one mile. Therefore, a cubic 
mile of such sand would contain 3 X io 6 cubed or 27 X 10 18 grains, 
which is the number of molecules in 1 c.c. of gas at standard 
pressure and temperature. Since the number of molecules per 
cubic centimeter has been determined by several independent 
methods which give closely agreeing results, we may safely 
accept the value given above as being correct within 1 per cent. 

196. Some Further Conclusions. — It may now be of some 
interest to note a few additional conclusions that have been 
reached in the study of gases. Let us illustrate by means of the 
gas oxygen. We know the weight of 1 c.c. of oxygen and the 
number of molecules in 1 c.c. at standard conditions; dividing 
the first by the second gives the weight of a single molecule of 
this gas; this comes out 5-3XIO -23 gram. The size of a mole- 
cule has also been approximately determined and in the case of 
oxygen it turns out that the diameter of a molecule is approxi- 
mately 2.5Xio~ 8 cm. We have already seen that the average 

distance between molecules is about of a cm. = 7. . 2 X 

3,000,000 

io -7 cm. We may now ask: How far, on the average, will a 
molecule travel in a straight line before it strikes another mole- 
cule? This result can be calculated when the diameter and 
average distance apart of the molecules are known, and is called 
the free path; for oxygen it is 1 . 3 X io -5 cm. Thus we see that, 
on the average, a molecule, after one collision will travel about 
40 times (i.3Xio _5 -t-3.3Xio _7 = 4o) the average distance 
between two neighboring molecules before striking a second 
molecule. This is not surprising when we note that the average 
distance between molecules is about 13 times their diameters 
(S'3X io" 7 ^- 2 . 5 X io~ 8 == 13). 

197. The Velocity of Molecular Motion. — The average ve- 
locity with which molecules travel between collisions can be calcu- 
lated with a high degree of certainty. The velocity varies with 
the mass of the molecule and its temperature but is independent 
of the pressure. Molecules of equal masses have equal velocities 



u6 Introduction to General Chemistry 

at the same temperatures, while for those with different masses 
the velocities are inversely proportional to the square root of the 
mass. At o° the velocity of the oxygen molecules is io 4 cm. per 
second, or about 15 miles per minute. But since the free path 
of a molecule of oxygen is only 1 . 3 X io~ 5 cm., it will experience 
many thousands of collisions in' progressing 1 cm. At each 
collision its direction of travel will change so that its actual 
progress from a given position is far slower than its high velocity 
would indicate if no collisions occurred. 

198. The Liquid State. — As a gas is compressed at constant 
temperature its molecules are brought closer together, but other- 
wise conditions remain nearly unchanged. The mass, diameter, 
and velocity of each molecule will not be altered; only the 
average distances between the molecules and their free paths 
will be shortened. It seems probable, in fact, that the average 
kinetic energy of a molecule, which is equal to one-half the 
product of its mass and the square of its velocity (§ niv 2 ), remains 
unchanged, however much the gas is compressed. If we accept 
this view, we may easily extend it to cover the liquid state, in 
which we may imagine that the molecules have the same veloci- 
ties and therefore the same kinetic energies as the molecules 
of the vapor of the liquid have at the same temperature, but that 
the crowding of the molecules is so great that their free paths 
are short compared with their diameters. However, we may 
think of the molecules as able to progress slowly from one place 
to another, although the motion will be very irregular, like that of 
persons moving about in a dense crowd. 

199. Vaporization of a Liquid. — It has already been stated 
that all of the molecules of a given gas cannot have equal veloci- 
ties; nor can a given molecule always have the same velocity, 
since at every one of the frequent collisions the velocity will be 
changed. It is only the average velocity of all the molecules 
that remains unchanged as long as the temperature remains 
constant. The velocities of the molecules of a liquid also are 
not all the same at a given instant; some will be moving much 
slower, others much faster, than the average. If a fast-moving 
molecule approaches the free surface of the liquid, it may escape 



Molecular Hypothesis 117 

into the space above the liquid, whereas a slow-moving molecule, 
under the same conditions, might not be able to escape. Now 
the passage of molecules from the liquid to the space above it is 
nothing but the evaporation of the liquid. Moreover, we see 
that the rate of escape of the molecules, and therefore the rate 
of evaporation, will be greater in proportion as the average 
velocity of the molecules is increased. Since molecular velocity 
increases with rise of temperature, we get in this way a simple 
explanation as to why heating a liquid hastens its evaporation. 
When the evaporation of a liquid goes on with a poor supply of 
heat, as when water evaporates in an open vessel, the liquid 
becomes cooler. Obviously this is due to the lowering of the 
average velocity of the molecules of the liquid because of the 
escape of the faster-moving ones. 

200. Vapor Pressure. — If a liquid is placed in a closed vessel 
which it does not completely fill, it will evaporate only until the 
pressure exerted by the vapor attains a certain value which is 
definitely determined by the temperature. For example, at 
20 the vapor pressure of water is equal to that exerted by 
17.4 cm. of mercury; at 25 it equals 23.6 cm. Does the water 
cease to pass into vapor when these pressures are reached ? If 
so, does this mean that molecules of water no longer pass from 
the liquid to the space above it? This would seem strange. 
Let us look at the question from another point of view. Suppose 
we have a vessel full of steam and allow it to cool.. We know 
that most of the steam will condense; only a little will remain as 
vapor. If we try to picture how this occurs, we must think of 
some of the molecules of vapor, that is, gaseous water, coming 
together first to form liquid droplets; these fall to the bottom 
and soon form a layer of liquid; other molecules then strike 
this liquid and remain as a part of it. Finally, when the tempera- 
ture of the room, say 20 , has been reached, the pressure within 
the vessel will have fallen to 17.4 cm. of mercury, and most of 
the water, but not all, will have condensed to the liquid state. 
It is important to note that at a given temperature, say 20 , the 
same final vapor pressure is reached whether steam condenses 
or water evaporates. 



n8 Introduction to General Chemistry 

201. Equilibrium between Liquid and Vapor. — A very impor- 
tant question now confronts us: Do water molecules cease to 
pass from the vapor into the liquid when at 20 the pressure 
reaches 17.4 cm.? If so, Why? Would it not seem more 
reasonable to suppose that/tfr every molecule that passes from the 
vapor into the liquid there is another that leaves the liquid and passes 
into the vapor ? This supposed state of affairs would correspond 
to that in which the number of customers in a large shop remains 
substantially constant during a given hour of the day, by reason 
of the fact that in each minute as many persons enter the shop 
as leave it. When at constant temperature the vapor pressure 
of a liquid has reached a constant value, we say there is equi- 
librium between liquid and vapor; and it would seem from the 
discussion above that this condition does not represent a state of 
rest or inaction, but one in which two opposing actions exactly 
counteract one another. 

202. Molecular Attraction. — If we think of the matter criti- 
cally, we may wonder why molecules of cooling water vapor col- 
lect into drops. Perhaps there is a sort of attraction between the 
molecules that holds them together. If so, why should it seem 
to be more effective at lower than at higher temperatures ? If, 
in reality, one molecule has some attraction for another, must 
we suppose that this attraction increases with fall of tempera- 
ture ? Would it not be sufficient to assume a constant attrac- 
tion of each molecule of water for every other? Suppose, now, 
the vapor of water is very hot ; then the molecules will be moving 
with such great velocities that if two of them collide they will 
rebound, exactly as a rubber ball, thrown downward, will 
rebound on striking the floor, although gravitational attraction 
tends to keep it on the floor. But suppose the vapor to be 
cooled; its molecules will then have smaller velocities and some 
may be moving so slowly that upon collision they remain in 
contact. Other slow-moving molecules, striking by accident 
a pair of molecules so formed, may add themselves to it, and 
in this way the droplet of water could be formed. There 
are also other reasons for assuming that molecules attract one 
another. 



Molecular Hypothesis 119 

203. The Solid State. — The most striking physical difference 
between a solid and a liquid is the rigidity of the former. This 
property of solids can most easily be accounted for by assuming 
that the molecules are not free to move about as in the case of a 
liquid, where the freedom of motion is comparable to that of 
people in a crowd, but that each molecule remains in its place 
with respect to the whole solid, as well as to its neighboring mole- 
cules. It is not necessary to think of the molecules as being 
absolutely at rest. It is more likely that each molecule has a 
vibrating motion at all temperatures above absolute zero and that, 
in fact, its kinetic energy is as great as it would be if the mole- 
cule were in the vapor state at the same temperature. 

204. Crystals. — Pure chemical substances in the solid state 
usually form crystals. The crystals of a given substance all have 
the same general form. Thus, for example, the crystals of 
common salt, when perfect, are all cubical in form, while those of 
quartz occur as hexagonal prisms. If we think of a crystal as 
built up of molecules, it is natural to wonder whether the mole- 
cules are present in haphazard fashion, like potatoes in a barrel, or 
if they may not perhaps be arranged in some systematic manner, 
like bricks in a wall or balls in a regular pile. The probability 
that the molecules of a crystal are arranged in a definite and regular 
manner is greatly increased when it is known that there are exactly 
as many types of crystalline form as there are possible regular 
arrangements of points in space. 

Within the last few years it has become possible by means 
of photographic studies made by the use of X-rays to obtain very 
precise information regarding the arrangement of molecules 
forming a crystal. As a result we now know quite definitely the 
molecular structure of a number of crystals. 

205. The Melting of Crystals. — Pure crystalline substances 
have definite melting temperatures; thus, ice melts at o°, and 
potassium nitrate at 339 . Increase of temperature must 
increase the intensity of molecular vibration; at some tempera- 
ture (the melting-point) this vibration seems to become so great 
that the systematic structure of the crystal is wrecked. Leaving 
only an irregularly mixed mass of molecules, forming the resulting 



120 Introduction to General Chemistry 

liquid. Crystals cannot be heated above their melting-points. 
Ice, for example, although it may be melting on the surface, is 
never hotter than zero. 

206. Supercooling. — On the other hand, water may be cooled 
2 or 3 degrees below zero without freezing, if it is kept quiet and 
is not in contact with ice. Such supercooled water immediately 
begins to freeze if touched with a piece of ice. This phenomenon 
is a common one and is easily explained. In order that the 
formation of a crystal can start, a certain minimum number 
of molecules must come together in the proper positions. But 
this exact arrangement of the several molecules necessary may 
not readily occur, especially as immediately above the freezing- 
point (which is the same as the melting-temperature) the mole- 
cules are vibrating so fast that they are just able to shake apart 
this regular arrangement (that is, to melt the crystal) . At a 
little lower temperature the molecular motion is less and therefore 
the conditions are more favorable for the starting of crystalliza- 
tion. However, if a crystal of the substance is present, then 
supercooling does not occur, but the liquid at once begins to 
crystallize (freeze) at the temperature of its melting-point. The 
reason is obvious: now each molecule that touches the crystal 
can find its proper lodging-place, and so crystalline growth can 
continue. 

207. Solutions. — In a solution the molecules of the dissolved 
substance must be very uniformly distributed; it would seem, 
therefore, that they may be moving about freely among the 
molecules of the solvent, being carried from place to place by 
their own motions. The process of dissolving of a substance 
would closely resemble that of evaporation, and the crystalliza- 
tion of a solid from its solution would correspond to the conden- 
sation of a vapor to a liquid. In fact, we may imagine that in the 
case of a saturated solution in contact with the crystals of a 
substance we have a state of equilibrium as a result of the passage 
of molecules into and out of the solution at exactly equal rates. 



CHAPTER XI 

THE ATOMIC HYPOTHESIS AND ATOMIC WEIGHTS 

208. Dalton's Atomic Hypothesis. — The application of the 
Atomic-Molecular Hypothesis to the explanation of chemical 
phenomena was first made by John Dalton of Manchester in 
1803. Long before this time Bernoulli had proposed the Kinetic- 
Molecular Hypothesis as an explanation of the physical behavior 
of gases, and Dalton, knowing this view of the nature of matter, 
sought to explain the difference in solubility in water of different 
gases as due to a possible difference in size of their molecules. 
But how could this imagined difference be discovered? At this 
date, 1803, the theory of the indestructibility of matter and the 
doctrine of elements were well established, owing to the work of 
Lavoisier, a quarter of a century earlier, as well as the labors of 
many able chemists of the intervening period. It was generally 
accepted that the formation of a substance was due to the union 
of the elements composing it and in many cases the proportions 
of the elements in a compound were already known — not very 
accurately, it is true, but at least approximately. Dalton wished 
to discover the relative weights of the ultimate particles of gases ; 
but in order to do this he would have to know, in the case of 
hydrogen and oxygen, for example, in addition to knowing the 
weight of oxygen that would combine with a given weight of 
hydrogen, the relative numbers of ultimate particles of the two 
gases that combine with one another in the formation of water. 
As Dalton had no experimental means of discovering the informa- 
tion he lacked, he simply assumed that one ultimate particle 
of hydrogen united with one ultimate particle of oxygen to give 
one ultimate particle of water, meaning by the expression ulti- 
mate particle essentially the same as the Greeks and later philos- 
ophers meant by the terms "atom" or "molecule," that is, the 
smallest possible particle of the substance. 

209. Finding the Relative Weights of Atoms. — Now, since 1 g. 
of hydrogen unites with 8 g. of oxygen to form 9 g. of water, 



122 Introduction to General Chemistry 

the ultimate particle or atom of oxygen must weigh eight times 
as much as the ultimate particle or atom of hydrogen; and the 
ultimate particle of water, in this case the molecule, must weigh 
nine times as much as an atom of hydrogen. In making such 
suppositions Dalton also assumed that all of the atoms of hydro- 
gen were exactly alike in size, weight, and all other properties; 
that each atom of oxygen was exactly like every other atom of 
this element, but entirely different from an atom of any other 
element. Dalton knew that the same pair of elements often 
form two or more compounds in which the constituents are 
present in different proportions. This forced him to assume 
also that in such cases the atoms unite, not only one to one, but 
also one to two, or one to three, etc. In order that the student 
may have a perfectly clear notion of the matter, we may sum- 
marize by stating that Dalton assumed that a molecule of water 
is composed of one atom of hydrogen and one atom of oxygen, 
and then reached the conclusion that an atom of oxygen was 
eight times, and a molecule of water nine times, as heavy as an 
atom of hydrogen. But Dalton did not know, as we can see 
clearly, whether one atom of hydrogen unites with one atom of 
oxygen or with two or three of oxygen, or whether two or perhaps 
three atoms of hydrogen unite with one of oxygen to form a 
molecule of water: it was all a guess. But it must also be clear 
that if we could discover the numbers of atoms of hydrogen and 
oxygen in a molecule of water we could find the relative weights of 
the two atoms, knowing the percentages of hydrogen and oxygen 
in water. Now the question is: How can we discover the num- 
ber of atoms of each kind in a molecule of a substance ? 

210. The Application of Avogadro's Hypothesis. — Suppose 
we accept Avogadro's suggestion that equal volumes of all gases 
at the same temperature and pressure contain the same number 
of molecules, and see to what conclusion we are led. Let us 
represent by N the number of molecules in 22.4 liters of any 
gas under standard conditions. Now according to Dalton's 
suggestion one molecule of a given substance will contain one, 
two, three, or some small whole number of atoms of a given 
element, but cannot, by reason of the assumed indivisible nature 



Atomic Hypothesis and Atomic Weights 123 

of an atom, contain a fraction of an atom. Let us consider the 
gas ammonia as an example. Ammonia is composed of 17.8 
per cent of hydrogen and 82.2 per cent of nitrogen, and nothing 
else. One molecule of ammonia, according to Dalton's sug- 
gestion, contains one, two, three, or four, or at least some small 
number of atoms of hydrogen. Now, if 22.4 liters of ammonia 
gas under standard conditions contain N molecules, then this 
volume of the gas must contain iXN, 2XN, $XN, or some 
small number of times N atoms. The least number of hydrogen 
atoms that could possibly be contained in 22 .4 liters of ammonia 
is N, but the true number may be 2 XN, which we may write 2N, 
or it may be greater, as 3N or 4^, so far as we know; only it 
must be N or some small whole number of times N if we assume 
that there are N molecules in 22.4 liters of the gas and also 
assume, with Dalton, that each molecule of the gas contains 
one, two, three, or some small number of atoms of hydrogen. In 
the case of any other gaseous compound of hydrogen we should 
conclude, according to Avogadro, that 22.4 liters of the gas 
contained N molecules and that each molecule contained one, 
two, three, or some other small number of hydrogen atoms, the 
smallest possible number being one atom of hydrogen to the 
molecule, and therefore that 22.4 liters of the gas would contain 
N, 2N, or 3N, etc., atoms of combined hydrogen. 

211. The Number of Atoms and Weight of Hydrogen in 22.4 
Liters. — According to Dalton all hydrogen atoms are alike and 
each has a definite weight, so that the weight of N atoms of 
hydrogen would be a perfectly definite weight of this element. 
The weight of 2N atoms of hydrogen would, of course, be twice 
that of N atoms, etc. It seems reasonable to think that it would 
be likely to happen that in some of the gaseous compounds of 
hydrogen the molecules would contain but one atom of hydrogen 
each. In such a case 22.4 liters would contain N atoms of 
combined hydrogen, having a definite weight. Now as such 
gases contain the minimum possible number of atoms of hydro- 
gen in each molecule, namely, one, and as we assume that all 
gases contain N molecules in 22.4 liters, then such gases would 
contain the minimum possible weight of hydrogen in this volume. 



124 Introduction to General Chemistry 

As a matter of fact we actually find that in 22 . 4 liters of the vari- 
ous gaseous compounds of hydrogen the weight of this element is 
in no case less than 1 g. In this volume of any definite gas there is 
either no combined hydrogen or there is at least 1 g.: the minimum 
weight of hydrogen is 1 g. 

212. The Explanation of the Laws of Minimum and Multiple 
Weights. — In other gaseous compounds of hydrogen we find in 
22.4 liters larger weights of combined hydrogen, but these 
weights are then either 2 g., 3 g., or some whole multiple of the 
minimum weight. Is it not logical then to think that in such 
gases as hydrogen chloride, where the minimum weight, 1 g., 
of combined hydrogen is contained in 22 . 4 liters of the compound 
gas, the molecule contains but one atom of hydrogen, and that 
in acetylene, where 2 g. of hydrogen are found in 22.4 liters 
each molecule contains two atoms of hydrogen, while in methane 
with 4 g. of hydrogen in the same volume, there are four atoms 
of hydrogen per molecule? Undoubtedly so. We see then 
that we have in the assumptions made by Avogadro and Dalton 
the basis of an explanation of the remarkable Laws of Minimum 
and Multiple Weights, which have been discovered by experi- 
ment; and, because of the agreement between theory and fact, 
we are inclined to think that perhaps the views of Avogadro and 
Dalton are correct. In any case we cannot fail to see that these 
hypotheses are useful, and that, indeed, is the criterion by which 
the worth of any hypothesis should be judged. 

213. Application of the Explanation to Other Elements. — The 
question now arises whether the simple explanations of the 
laws of minimum and multiple weights may be applied to ele- 
ments other than hydrogen, and a very little thought will show 
that this must be "the case. The minimum number of atoms 
of any given element in 22.4 liters of any of its gaseous com- 
pounds is again N, the number found in the case of those gases 
the molecules of which contain but one atom each of the given 
element. The minimum weight of this element is, of course, the 
weight of N atoms of the element. For example, we find that 
in 22.4 liters of gaseous carbon compounds the minimum weight 
of carbon is 12 g. Since this minimum weight is found in the 



Atomic Hypothesis and Atomic Weights 125 

gases carbon dioxide and methane, we conclude that but one 
atom of carbon is contained in a molecule of each. On the other 
hand, 22.4 liters of acetylene contain 24 g. of combined carbon, 
which is twice the minimum weight, from which we conclude 
that in a molecule of this gas there are two atoms of carbon. 
As we have now reached the conclusion that a molecule of methane 
contains four atoms of hydrogen and one of carbon we see that we 
have developed a method whereby we can solve the problem 
first suggested by Dal ton, that of discovering the number of atoms 
of each sort in a molecule of a given substance at least in the case 
where the substance is a gas, since it is evident that the method 
used for methane is applicable to any gaseous substance. 

214. The Number of Atoms of Each Kind in a Molecule. — To 
illustrate by further examples, we may consider the cases of a 
few of the gases we have already studied. We see, by reference 
to Table IV, that in 22.4 liters of hydrogen chloride there are 
found the minimum weights of both hydrogen and chlorine and 
conclude that the molecule of this gas is made up of one atom 
each of hydrogen and chlorine. We also see by Table IV that 
the unit volume of ammonia contains the minimum weight of 
nitrogen and three times the minimum weight of hydrogen, and 
decide that in a molecule of ammonia one atom of nitrogen must 
be united with three atoms of hydrogen. In all other cases the 
reasoning is equally simple, so that the student will have no 
trouble in deciding upon the number of atoms of each kind in a 
molecule of each of the gases mentioned in Table IV. 

215. The Number and Kind of Atoms in a Molecule Shown 
by the Formula. — We are now in position to notice a most re- 
markable fact, which the following examples will illustrate. One 
molecule of hydrogen chloride contains one atom of hydrogen and 
one of chlorine , and its formula is HCl; one molecule of ammonia 
contains three atoms of hydrogen and one atom of nitrogen, and 
its formula is NH 3 ; one molecule of methane contains four atoms 
of hydrogen and one atom of carbon, and its formula is CH 4 ; 
in each case the number of atoms of each element is the same as the 
number of symbol weights of that element in the formula of the sub- 
stance! And that the same thing is true for all gases of Table IV 



126 Introduction to General Chemistry 

may readily be found by considering each separate case in the 
same way as we did those of three of the gases. In every case, 
therefore, the formula shows not only the weight of each element in 
22 . 4 liters of the gas but also the number of atoms of each element 
in one molecule of the substance. 

216. Symbol Weights and Atomic Weights. — But we may 
now inquire, Why should this be true? To answer this question, 
we will recall that the minimum weight of any element in 22.4 
liters of its gaseous compounds is the weight of X atoms of that 
element. If N atoms of hydrogen weigh 1 g. and N atoms of 
carbon weigh 12 g., then one atom of carbon must be 12 times as 
heavy as an atom of hydrogen. In a similar way we are led to 
conclude that an -atom of nitrogen is 14 times, and an atom of 
oxygen 16 times, as heavy as an atom of hydrogen. Analogous 
relations must likewise exist in the cases of all other elements; 
and, therefore, taking the weight of one atom of hydrogen as one 
or unity, the weight of an atom of any other element is repre- 
sented by exactly the same number as its symbol weight. For 
this reason a table of symbol weights is also called a table of 
Atomic Weights ; and symbol weights are usually referred to as 
atomic weights. But we must remember that the symbol 
weights may be found by simple and direct experiments, inde- 
pendently of all suppositions and hypotheses, while atomic 
weights are to be represented by the same set of numbers only 
when we assume that matter is made up of atoms which unite 
in simple ratios to form molecules of which all gases are assumed 
to contain equal numbers in equal volumes. Briefly stated, 
symbol weights are natural constants, but atomic weights are the 
probable relative weights of the atoms of which we imagine matter 
to be made up. We now may answer the question proposed in the 
first sentence of this paragraph. The number of atoms of any 
sort in a molecule is the same as the number of symbol weights of 
that element because the absolute weight of an atom of any element is 
proportional to its symbol weight. In this chapter we have seen 
how the problem which Dalton set for himself over a century 
ago is to be solved, at least as definitely as chemists know, up 
to the present time, how to solve it. The key to the solution was 



Atomic Hypothesis and Atomic Weights 127 

the hypothesis of Avogadro, which was suggested in 181 1, only 
three years after Dalton's views first appeared in print, and 
which was rejected by Dalton himself, and was only accepted 
by the chemical world at large half a century later. 

217. Formula Weights and Molecular Weights. — It is of 
course obvious that the weight of a molecule may also be 
expressed in terms of the weight of one atom of hydrogen which 
is taken as unity. For example, if one atom each of hydrogen 
and chlorine compose a molecule of hydrogen chloride and if, as 
we have seen, an atom of chlorine weighs 35. 5 times as much as 
an atom of hydrogen, then a molecule of the compound must 
weigh 36. 5 times as much as an atom of hydrogen; and we say 
therefore that the Molecular Weight of hydrogen chloride is 
36.5. The molecular weight of a gas has consequently the same 
numerical value as its formula weight, the weight of an atom of 
hydrogen being in all cases taken as unity. The conclusion 
that the relative weights of the molecules of gases are propor- 
tional to their respective formula weights follows at once from the 
assumption of Avogadro's hypothesis. But we now see also that 
the weights of gaseous molecules, briefly their molecular weights, 
are all represented by the same numbers as their formula weights 
if we choose the atomic weight of hydrogen as unity; for this 
reason it seemed logical to discuss atomic weights before molec- 
ular weights. // is also evident that the molecular weight of a 
substance must be equal to the sum of the atomic weights indicated 
by the formula. 

218. The Formulae of Some Elementary Gases.— We are now 
in position to consider the meaning of the fact that the formulae 
H 2 , 2 , N 2 , and Cl 2 were found for the four elementary gases 
studied. If we accept Avogadro's hypothesis for these as well 
as for compound gases, then the unit volume of each gas must 
also contain N molecules. But we know also that the weight 
in each case is twice the minimum weight of N atoms; for this 
reason we are forced to conclude that the unit volume of each 
gas contains 2 A 7 atoms, and hence that each molecule contains 
two atoms. The same rule that applies to compound gases 
applies here also: the number of atoms of each element in any 



128 Introduction to General Chemistry 

molecule is the same as the number of symbol weights of that 
element. The obvious meaning of these facts is that an atom 
of hydrogen, for example, can unite with another atom of hydro- 
gen as well as with one of chlorine or some other element. But 
if this is the conclusion, you will doubtless ask : Do we have any 
other evidence of its truth? Let us see. 

219. The Union of Hydrogen and Chlorine by Volume. — It 
will be recalled that when hydrogen and chlorine gases unite to 
form gaseous hydrogen chloride, one volume of each of the 
elementary gases combines to give two volumes of the product. 
Now in two unit volumes of 22 . 4 liters each of hydrogen chloride 
there must be 2 N atoms of hydrogen and 2N atoms of chlorine, 
since we cannot have less than one atom of each element in a 
molecule of the compound ; but the two unit volumes of hydrogen 
chloride are formed from one unit volume of hydrogen and one of 
chlorine each containing N molecules; and again we are led to 
the conclusion that N molecules of hydrogen or chlorine contain 
2 N atoms in each case, or that each molecule of either elemen- 
tary gas contains two atoms. 

220. Gay Lussac's Law of Combining Volumes. — The simple 
volumetric relation between hydrogen, chlorine, and hydrogen 
chloride, 1:1:2, is not an exceptional case; other gases also 
exhibit similar simple relations. Thus, two volumes of hydrogen 
and one volume of oxygen unite and, if the temperature at which 
the experiment is carried out is so high that the water remains in 
the form of steam, the latter measures two volumes; so that the 
volume relations are 2:1:2. In the burning of methane one 
volume of the gas requires two volumes of oxygen and gives 
one volume of carbon dioxide and two volumes of steam, the 
measurements all being made at a sufficiently high temperature 
in this case also to keep the steam from condensing. Or, when 
ammonia gas is decomposed, as it may be by means of electric 
sparks, two volumes of ammonia yield one volume of nitrogen 
and three volumes of hydrogen. The fact that gases and vapors 
of volatile substances always react in simple ratios by volume 
was discovered by Gay Lussac in 1808, and is known as Gay 
Lussac's Law of Combining Volumes. 



Atomic Hypothesis and Atomic Weights 129 

221. Explanation of the Law of Combining Volumes. — The 

explanation of this law will appear if we write the equations of 
the reactions mentioned: 

H 2 +C1 2 ^ 2 HC1 CH 4 + 2 2 ->C0 2 +2H 2 

I Vol. I Vol. 2 Vols. I VOL 2 VOls. I Vol. 2 Vols. 

2H 2 +0 2 ->2H 2 2 NH 3 -> 3 H 2 +N 2 

2 VOls. I Vol. 2 VOls. 2 Vols. 3 Vols. I Vol. 

Again we see, as we did earlier (76, 77), that the volumes are 
the same as the coefficients of the formulae in the equations, and 
this for the fundamental reason that one formula weight always 
represents one unit volume, in the case of a gaseous or volatile 
substance. Moreover, we now understand, to cite the first 
example, that one unit volume of hydrogen containing N mole- 
cules and 2 N atoms will require 2 N atoms of chlorine or N mole- 
cules, which, according to Avogadro's hypothesis, will be found 
in one unit volume of chlorine. The reaction will then produce 
2 N molecules of hydrogen chloride, which according to the same 
hypothesis will occupy two unit volumes. Similar reasoning 
may be applied to all other cases. 

222. The Degree of Accuracy of Symbol Weights. — Before 
leaving the discussion of symbol and atomic weights we must 
consider the degree of accuracy of the statements of numerical 
results made in early chapters and summarized in Table IV. It 
is perhaps needless to point out that statements of lengths, areas, 
volumes, weights, etc., whether they refer to scientific or other 
matters, are in general more or less approximate, the degree of 
accuracy aimed at being determined by the requirements of the 
case. Thus, if a stranger in the city inquires the distance from 
the City Hall to the University and is told by a policeman that 
it is seven miles, the answer is quite as accurate as necessary. 
But such approximate statements of distance would not satisfy 
the requirements of a surveyor who wished to make an accurate 
map of the city. Up to about twenty-five years ago the most 
accurate analyses of water indicated that 2 g. of hydrogen were 
combined with 15.96 g. of oxygen. As all chemists know that 
in every analysis there is inevitably some experimental error 



130 Introduction to General Chemistry 

of greater or less magnitude, it was thought that the true weight 
of oxygen combined with exactly 2 g. of hydrogen was exactly 
16 g. It then became apparent from the new researches of a 
number of chemists that the error in the accepted results was 
greater than suspected, and, moreover, that the true proportion 
of oxygen in water was less instead of greater than the value 
found earlier, the new experiments leading to a ratio of 2 to 15 . 88, 
with a probable error of less than 0.01 g. in the weight of oxygen 
combined with exactly 2 g. of hydrogen. 

223. = i6.ooo, the Real Basis for Symbol and Atomic 
Weights. — An annoying difficulty now arose from the fact that 
far more symbol weights had been found by the analysis of oxygen 
compounds than by the analysis of compounds with hydrogen, 
owing to the greater accuracy with which the former analyses 
could be made; so that it then became necessary for chemists to 
decide whether they should change the symbol weights of oxygen 
and all elements whose symbol weights had been found by the 
analysis of their oxygen compounds, or whether they should 
change the symbol weights of hydrogen and a few other elements. 
After much debate the former policy was adopted and the symbol 
weight of oxygen, = 16.000, kept unchanged, although this 
made it necessary to change the symbol weight of hydrogen to 
1.008. Our most accurate knowledge of the composition of 
water is expressed by the statement that 2X1. 008 g. of hydrogen 
are combined with 16.000 g. of oxygen in 18.016 g. of water, a 
fact which is also expressed by the formula H 2 0, when we consider 
that H = 1.008 g. of hydrogen and = 16.000 g. of oxygen. 
Oxygen with a symbol weight of 16.000 has thus become the real 
basis of the system of symbol and atomic weights rather than hydrogen 
with a symbol weight of unity. 

224. The Method of Finding Symbol Weights. — The symbol 
weights, and therefore also the atomic weights, of all other ele- 
ments are now based upon that of oxygen taken as 16.000; but 
we see by a comparison of the values given in a table of exact 
atomic weights that in no case does the exact value based on 
= i6.ooo differ greatly from the approximate value we have 
previously used. Just as more accurate analyses led to a change 



Atomic Hypothesis and Atomic Weights 131 

in the symbol weight of hydrogen, so also newer analyses have 
led and will continue in the future to lead to a more exact 
knowledge of the symbol weights of other elements. We do not 
expect, however, that the values accepted at present for the 
commoner elements will be changed by more than a few units 
in the second decimal place. Concisely stated, the matter stands 
thus: Approximate symbol weights are found in the manner 
described in chap, v, while the more exact values are fixed by the 
most painstaking analyses and syntheses, being computed on the 
basis of = i6.ooo. 

225. Inexactness of the Gas Laws. — The gas laws of Boyle, 
Charles, and Avogadro are also only closely approximate state- 
ments of the facts. For example, if the pressure on 1,000 c.c. 
of oxygen under standard conditions be exactly doubled, the 
volume will become 499.3 c.c. instead of exactly 500, as Boyle's 
law would indicate. The deviations from the simple laws are 
thought to be due to attractions between the molecules, on the 
one hand, tending to diminish the volume, and, on the other 
hand, to the fact that part of the space occupied by the gas is 
rilled with the molecules themselves, so that the, free space is 
reduced to less than half if the volume of the gas is, by increase 
of pressure, reduced to half. The actual deviations from the 
simple law, PV = sl constant, become negligible if gases are under 
low pressures. Then the three great laws express almost exactly 
the behavior of all gases. In other words, if the barometric 
pressure at sea-level were 0.01 of its actual value, so that our 
standard of atmospheric pressure would be o. 76 cm. of mercury 
instead of 76 cm., then we should find that not only would the 
laws of Boyle and Charles express with a high degree of accuracy 
the behavior of gases under pressures of this order of magnitude, 
but that for all gases the law of Avogadro would also hold good 
with as great a degree of accuracy as experiment would enable 
us to determine. 

226. Exactness of Avogadro' s Law for Corrected Gas Vol- 
umes. — Now, instead of trying to weigh and measure gases under 
such low pressures in attempts to study them more accurately, 
chemists have worked at ordinary pressures and then corrected 



132 Introduction to General Chemistry 

the data so obtained so as to give the results that would theoreti- 
cally have been obtained for the weights of 1 liter if the measure- 
ments had been made at very low pressures and the calculations 
made for a pressure of 76 cm. exactly according to Boyle's law. 
Working in this way, it was found that the corrected volume of 
32 g. of oxygen, the weight represented by 2 , is 22.41 liters 
at o°. It was then discovered that exactly this (corrected) 
volume of any other gas at o° contains, as nearly as the deter- 
minations could be made, just the weight of the gas which its 
formula indicates, this weight being calculated from the most 
exact symbol weights. In others, Avogadro's law would hold 
exactly at low pressures or also at ordinary pressure if the attrac- 
tions of the molecules for each other did not exist, and if their 
own volumes were negligible as compared with the total space 
occupied by the gas. 

227. A Little Explanation and Advice. — It is not necessary nor 
desirable that the beginner in chemistry should pay much atten- 
tion to the matters discussed in the three preceding paragraphs. 
The approximate symbol weights and the gas laws in their sim- 
plest forms are sufficiently exact for his use. It is much better 
that he should see clearly the general fundamental principles 
than that he should be perplexed and confused by the details 
and refinements that are of importance only to the specialist. 
If the beginner continues his study of chemistry he will be sure 
to encounter later these interesting topics, when he will be better 
able to appreciate and understand them; while if he should not go 
farther than the first course, he may feel assured that he has be- 
come acquainted with the principles of most fundamental impor- 
tance. These matters are discussed here in order to explain why 
the symbol or atomic weights given in Tables of Atomic Weights, 
(see inside of back cover of this book) are not exactly the same 
as those we have used in the earlier chapters. 

228. Means of Discovering Symbol Weights. — The student 
will doubtless have received the impression from the study of the 
foregoing chapters that we can discover the approximate symbol 
weight of any element by finding the minimum weight of the 
element in the unit volume of its gaseous or vaporized com- 



Atomic Hypothesis and Atomic Weights 



*33 



pounds; and this, in fact, is true for a large number of elements 
in addition to the five included in Table IV. We shall now 
consider some facts leading to a knowledge of the symbol weights 
of a dozen elements other than the five already studied. These 
twelve elements all form volatile compounds, the densities of 
which may be determined by making experiments at sufficiently 
high temperatures and then calculating, by the laws of Boyle 
and Charles, for the standard, conditions, the weight of the com- 
pound in 22.4 liters. Multiplication of the weight so found by 
the percentage of the element in question in the compound gives 
the weight of the element in 22.4 liters of the vapor, as recorded 
in Table X. 

TABLE x 



Volatile Compounds of 
Various Elements 

Antimony trichloride . . 
Arsenic trichloride. . . . 
Bismuth trichloride. . . 

Cadmium 

Chromium oxychloride 

Hydrogen iodide 

Iron carbonyl 

Lead chloride 

Mercury 

Nickel carbonyl 

Phosphorus trichloride 
Zinc chloride 



Weight of 
Elements in 
22.4 Liters 



Symbol and 
Symbol Weight 



Specific Heat 



Product of 
Symbol Weight 
and Specific Heat 



IIQ 

75 
217 
114 

55 
127.7 

53-2 
207 
202 

59 

3i 

63 



Sb =120.2 
As = 75.0 
Bi =208.5 
Cd =112.4 
Cr = 52. 1 
I =127.9 
Fe = 55.9 
Pb = 206 . 9 
Hg = 200 . o 
Ni = 58.7 
P = 310 
Zn = 65.4 



0.0503 
o . 0830 
0.0303 

0.0551 

O. II2I 
O.054I 

o. 1162 
o . 0304 
o . 0308 
0.1084 
o. 2020 

0.0935 



6.0 

6.2 

6.3 

6.2 

5-8 
6.9 
6-5 
6-3 
6.1 
6.4 
6-3 
6.1 



In all cases the compounds are such as contain the mini- 
mum weight of the element the symbol of which appears in the 
table; that is to say, we do not know any other volatile com- 
pounds in the respective cases containing appreciably smaller 
weights in the unit volume. The weights so found are, therefore, 
approximately the symbol weights in each case. The exact 
symbol weight in any case is then calculated from the accurately 
determined percentage composition of some compound of the 
element with an element of exactly known symbol weight. 

229. The Product of Specific Heats and Symbol Weights. — 
There are a great many elements which do not form gaseous com- 
pounds, or compounds which are sufficiently volatile without 



134 Introduction to General Chemistry 

decomposition, to enable us to find their symbol weights in the 
manner above indicated. Very fortunately other methods have 
long been known by which the desired end can be attained. We 
shall now consider one of these methods. 

A very simple relation was discovered nearly a century 
a g°> by Dulong and Petit, between symbol weights and specific 
heats of solid free elements. The amount of heat required to 
raise the temperature of a given weight of iron i° would raise- 
the temperature of an equal weight of water only o. 1162 ; and 
we say, therefore, that the specific heat of iron is o. 1162. The 
specific heats of the other elements of Table X are given in the 
fourth column. If, now, we multiply the specific heat of an 
element by its symbol weight we get the remarkable series of 
products contained in the last column of the table, where we see 
that the values are nearly the same in all cases. Does it not 
seem probable that the law which we find applying to the ele- 
ments of Table X would also hold good for other solid elements 
even though they do not form easily volatile compounds? If 
so, it is clear that in order to find the approximate symbol weight 
of an element we have only to divide 6.4 by its specific heat, which 
latter constant can in general be found by a simple, direct 
experiment. As a matter of fact, this method has been of much 
service in just this way. 

230. Interpretation of the Law of Dulong and Petit. — The law 
of Dulong and Petit is, moreover, of the greatest interest and 
importance when viewed from the theoretical standpoint. The 
product of the specific heat and symbol weight is obviously the 
quantity of heat required to raise the temperature of the symbol 
weight of an element one degree ; and this amount of heat is the 
same for one element as for another. But the symbol weights 
of various elements are the weights of equal numbers of atoms, 
and we see, therefore, that it requires equal amounts of heat 
to raise the temperature of equal numbers of various kinds of 
atoms by one degree! The products of symbol weights and 
specific heats are generally called Atomic Heats; so that the 
Law of Dulong and Petit may be stated thus: The atomic heats 
of the solid elements are equal. 



CHAPTER XII 

THE HALOGENS AND THEIR COMPOUNDS WITH HYDROGEN 

AND METALS 

231. The Halogens. — The elements fluorine, chlorine, bro- 
mine, and iodine bear a close resemblance to one another in their 
properties and chemical behavior; collectively they are called 
the halogens (from halite, the scientific name for rock salt). 
In the present chapter we shall first briefly review what has 
already been learned about chlorine and some of its compounds, 
and then after a more extensive consideration of the chemistry 
of chlorine take up a study of the remaining members of this 
important group of elements. 

232. Resume of Facts Already Learned. — We know that 
common salt, NaCl, is the most abundant compound of chlorine; 
it forms the raw material from which all other compounds of 
chlorine as well as the free element are made. The action of 
sulfuric acid on salt (103) yields hydrochloric acid which, 
by electrolysis (43) or by the action of lead dioxide, gives free 
chlorine (167). With bases or metallic oxides hydrochloric acid 
yields chlorides, as illustrated by the following reactions: 

K0H+HC1->KC1+H 2 (107) 

MgO+2HCl->MgCl 2 +H 2 0. (143) 

Chlorides also result when carbonates are treated with hydro- 
chloric acid (163) : 

Na 2 C0 3 + 2HC1^2NaCl+ C0 2 +H 2 
CaC0 3 +2HCl^CaCl 2 +C0 2 +H 2 0. 

It will be recalled that the chlorides of silver, lead, and 
univalent mercury are almost insoluble in water (167, 169, 182); 
these salts are easily obtained by the action of solutions of hydro- 
chloric acid or any soluble chloride on solutions of soluble salts 
of these metals, thus: 

AgN0 3 +HCl^AgCl+HN0 3 

Pb(N0 3 ) 2 +2NaCl->PbCL+2NaN0 3 . 

135 



136. Introduction to General Chemistry 

The metals which react with hydrochloric acid set free hydrogen 
and are themselves converted into chlorides, for example: 

Zn+ 2 HCl^ZnCl 2 +H 2 (149) 

2 A1+6HC1->2A1C1 3 +3H 2 . (174) 

Chlorides also result from the direct union of chlorine with other 

elements: 

H 2 +C1 2 ->2HC1 (44) 

2Al+ 3 a->2AlCl 3 . (174) 

233. The Occurrence of Chlorine Compounds in Nature. — 
Free chlorine does not occur in nature. If free chlorine were 
present in nature it would very soon unite with other substances 
to form compounds. Common salt is by far the most abundant 
natural compound of the element. It occurs as a mineral, rock 
salt (halite) , and as dissolved salt in sea-water and the waters of 
salt lakes and springs. Sea-water contains about 3 per cent, 
while the water of Great Salt Lake in Utah contains about 
20 per cent, of salt. Rock salt has doubtless been formed in 
past geological times by the slow, natural evaporation of sea- 
water. Other chlorides, particularly those of potassium, KC1; 
magnesium, MgCl 2 ; silver, AgCl; and lead, PbCl 2 , are also found 
in nature. 

234. The Discovery of Chlorine. — Free chlorine was first 
made by the Swedish chemist Scheele, in 1774, and therefore 
practically at the same time that Lavoisier in France discovered 
the true explanation of burning. Scheele made chlorine by 
the action of hydrochloric acid on manganese dioxide, a mineral 
having the formula Mn0 2 , and therefore an oxide of the metallic 
element manganese. The reaction occurs thus: 

4 HCl+Mn0 2 ->MnCl 2 +Cl 2 +2H 2 0. 

Chlorine was not thought to be an element until nearly forty 
years after its discovery, but was believed to be an oxide of 
hydrochloric acid, until a famous English chemist, Sir Humphrey 
Davy, showed by conclusive experiments that it did not contain 
oxygen and was really an elementary substance. 



Halogens with Hydrogen and Metals 



137 



235. The Preparation of Chlorine from Hydrochloric Acid. — 
We have already seen (167) that chlorine is formed when lead 
dioxide is warmed with hydrochloric acid : 

4 HCl+Pb0 2 ->PbCl 2 +Cl 2 +2H 2 0. 

This reaction is entirely analogous to the one between hydro- 
chloric acid and manganese dioxide mentioned in the preceding 
paragraph, and since the last substance is cheaper than lead 
dioxide it is the one commonly used in the laboratory for the 
preparation of chlorine. The experimental method consists in 
adding to, say, 100 g. of granular manganese dioxide contained 
in a flask about 300 c.c. of concentrated 
hydrochloric acid and warming gently: 

4 HCl+Mn0 2 ->MnCl 2 +Cl 2 +2H 2 0. 

Manganese chloride is an easily soluble 
salt which forms pink crystals of a 
hydrate, MnCl 2 • 4H 2 0. 

An excellent, though expensive, 
method of making small amounts of 
chlorine for experimental work in the 
laboratory consists in allowing concentrated hydrochloric acid to 
drop slowly onto solid potassium permanganate, KMn0 4 (Fig. 30). 
The latter substance is one of the most powerful oxidizing agents 
and reacts rapidly in the cold with hydrochloric acid, thus : 

i6HCl+2KMn0 4 ->2KCl+2MnCl 2 +5Cl 2 +8H 2 0. 

Since the rate of production of chlorine is easily regulated by 
control of the rate of flow of the acid, the method is a very con- 
venient one for the lecture table. 

236. Chlorine, a Poisonous Gas. — The chlorine which is given 
off is a heavy, yellowish, poisonous gas having an exceedingly 
violent action on all mucous membranes. It is the gas which 
was first used with such frightful effect in the trenches in the 
European war. Great care must be exercised to prevent the escape 
of appreciable amounts of chlorine into the air of the laboratory 




Fig. 30 



138 Introduction to .General Chemistry 

and to avoid as far as possible inhalation of the gas. Waste 
chlorine is easily absorbed when passed into a solution of 
caustic soda. 

237. The Electrolytic Preparation of Chlorine. — We have 
already learned (43) that chlorine is formed when hydrochloric 
acid is electrolyzed. By means of the Brownlee apparatus 
shown in Fig. 21 it is found that equal volumes of hydrogen 
and chlorine are formed when the concentrated acid is used. 
If, however, very dilute acid is used, then the products are 
largely hydrogen and oxygen formed by the decomposition of 
the water, and very little chlorine is set free. A complete 
explanation of this curious fact is not possible until certain 
matters treated in a following chapter have been considered; 
but it may be stated that hydrochloric acid is more easily 
decomposed than water by the electric current, and that if much 
of the former is present in a water solution it is decomposed by 
preference to the water. In the electrolysis apparatus the poles 
or electrodes are sticks of carbon. The hydrogen is liberated 
at the negative pole, the chlorine at the positive pole. 

238. The Electrolysis of Common Salt. — The electrolysis of 
a concentrated solution of common salt is by far the most im- 
portant practical method for the manufacture of chlorine. It 
is a process which is carried out on a very large scale, as at 
Niagara Falls, where electrical power is cheap and yields not only 
chlorine but also hydrogen and caustic soda. We might expect 
the products of the electrolysis of salt to be sodium and chlorine, 

2 NaCl->2Na+Cl 2 , 

but when we recall that sodium reacts at once with water to 
form hydrogen and sodium hydroxide (caustic soda), the actual 
result appears reasonable. A more complete explanation must 
be deferred until later. As in the case of the electrolysis of 
hydrochloric acid the chlorine is set free at the positive electrode, 
which is a carbon plate, while the sodium and hydrogen are 
formed at the negative electrode. 

239. Deacon's Process. — Before the electrical method just 
described was used practically, a process invented by Deacon 



Halogens with Hydrogen and Metals 139 

was the cheapest technical method of making chlorine. This 

process is based on the fact that a mixture of hydrogen chloride 

gas and oxygen react at a high temperature to form chlorine and 

water, 

4 HCH-0 2 ^2C1 2 +2H 2 0. 

This reaction scarcely takes place at all at ordinary temperatures, 
and even at the most favorable high temperature it takes place 
very slowly. Deacon discovered that the reaction could be 
greatly hastened if the heated mixture of hydrogen chloride and 
oxygen were passed over broken bricks coated with copper 
chloride, CuCl 2 . A small amount of this substance is able to 
promote the reaction of almost unlimited amounts of the reacting 
gases without itself being permanently changed or destroyed. 
A substance that behaves in this way is called a catalytic agent. 
Catalytic agents of various sorts are extensively employed in 
chemistry. In the Deacon process air, which is essentially a 
mixture of oxygen and nitrogen, may be used instead of pure 
oxygen, which would be too expensive for practical purposes. 

240. A Remarkable Phenomenon : Chemical Equilibrium. — 
It is a remarkable fact that even under the most favorable con- 
ditions the reaction between hydrogen chloride and oxygen does 
not go to completion, but stops while the gaseous mixture still 
contains some of both of these gases. The cause is discovered 
when we find that steam and chlorine react at about 400 to give 
some hydrogen chloride and oxygen: 

2Cl 2 +2H 2 0-> 4 HCl+0 2 . 

This is, in fact, exactly the reverse of the reaction we have been con- 
sidering. It is plain, therefore, that the failure of the reaction 
between hydrogen chloride and oxygen to go to completion is 
due to the interaction of the products, chlorine and water, to 
form again some of the first pair of gases. 

If a mixture of hydrogen chloride and oxygen in the propor- 
tions shown in the equation is heated to a constant temperature, 
say 400 , a mixture finally results in which all four of the sub- 
stances are present in definite proportions. A mixture having 
exactly the same proportions of each of the four substances 



140 



Introduction to General Chemistry 




Fig. 31 



results if the starting substances are chlorine and water, taken 
also in the proportions indicated by the equation. In the 
mixture which finally results, the four substances are said to be 
in a state of chemical equilibrium. The subject of chemical 
equilibrium is a very important one which is to be studied in 
detail in the next chapter. 

241. The Physical Properties of Chlorine. — Chlorine is a 
pale-yellow gas, having a density about two and a half times as 
great as air. Under standard conditions one liter weighs 3 . 22 g. 
Chlorine is rather soluble in water, 100 c.c. of water at 20 dis- 
solving 226 c.c. of the gas. For this reason the gas is not easily 

collected over water ; on account 
of its high density it is easily 
collected by the downward dis- 
placement of air. If a water 
solution of chlorine is cooled 
nearly to o°, yellow crystalline 
chlorine hydrate, having the 
formula C1 2 -8H 2 0, is formed. 
This hydrate is very unstable and decomposes slowly at room 
temperature and rapidly at higher temperatures into chlorine 
gas and water. 

242. The Liquefaction of Chlorine. — A very interesting and 
important experiment was once made with this hydrate by the 
great English physicist and chemist Faraday, who was at the 
time assistant to Sir Humphrey Davy (234). Crystals of 
chlorine hydrate were sealed up in one end of a bent glass tube, 
as shown in Fig. 31; when the hydrate was gently warmed 
while the other end of the tube was cooled with ice a yellow 
liquid formed in the cold end of the tube. This liquid proved 
to be liquefied chlorine. It is a heavy, mobile liquid, which is 
easily obtained from chlorine gas either by cooling the latter 
to about 40 below zero at atmospheric pressure, or by com- 
pressing it to about four atmospheres' pressure at about o°. 
Under one-atmosphere pressure liquid chlorine boils at — 34 . 
This work of Faraday in liquefying chlorine was of very great 
importance, since it was the -beginning of the epoch-making 



Halogens with Hydrogen and Metals 



141 



experiments in which he succeeded in liquefying all known 
gases except rive, among which were hydrogen, oxygen, and 
nitrogen. 

243. The Union of Chlorine and Hydrogen. — Chlorine and 
hydrogen do not react at an appreciable rate at room tempera- 
ture if kept in complete darkness, but do unite with explosive 
violence if exposed to a bright light, hydrogen chloride being 

formed, thus: 

H 2 +C1 2 ->2HC1. 

In order to demonstrate this interesting phenomenon a thin- 
walled glass bulb is filled with a mixture of equal volumes of 
the two gases; the bulb is then covered with a thick- walled bell 
jar (Fig. 32) and strongly illuminated either by direct sunlight 




m 




tt 



m 



Fig. 32 



Fig. 33 



or by the rays from burning magnesium ribbon. The sharp 
explosion which follows reduces the glass bulb to a powder, but 
does no damage to the bell jar. The mixture of chlorine and 
hydrogen is best obtained by the electrolysis of concentrated 
hydrochloric acid in the apparatus shown in Fig. 33. The inner 
vessel has two carbon electrodes. It is surrounded by a larger 
vessel, through which water flows to prevent rise of temperature. 
During the filling of the bulb and up to the time all is ready for 
the explosion it must be shielded from bright light. The union of 
chlorine with hydrogen takes place slowly, without explosion, if 
the mixture of the two gases is exposed for a sufficient length of 
time to moderate light (44). 

244. The Burning of Hydrogen in Chlorine. — If a jet of hydro- 
gen burning in air is lowered into a jar of chlorine it continues 



142 



Introduction to General Chemistry 



r 



to burn with a pale flame (Fig. 34). The flame is the result of 

the intense heat produced by the union of the two gases to form 

hydrogen chloride. 

245. The Action of Chlorine on Water. — Water dissolves 

about two or three times its own volume of chlorine at room 

temperature, giving a yellowish solution 
known as chlorine water. This solution 
smells strongly of chlorine and is often 
used in the laboratory in place of 
chlorine gas. If chlorine water is 
exposed to light it soon loses its color 
and odor, and at the same time a color- 
less, odorless gas, which proves to be 
oxygen, is given off. The experiment 
may readily be carried out in the 
manner shown in Fig. 35. A cylinder 
filled with chlorine water is inverted in 

a dish or beaker and exposed to bright light for a day or two. 

The gas produced will be found to be oxygen, formed according 

to the equation 

2Cl 2 + 2 H 2 0^ 4 HCl+0 2 . 



Fig. 34 



This is the reversal of the reaction by which chlorine is made by 
Deacon's process. While chlorine gas and steam react only 
partially at a high temperature, as already 
stated, chlorine dissolved in water and ex- 
posed to light reacts slowly, but completely, at 
room temperature to form hydrochloric acid 
and oxygen. This curious difference in be- 
havior may be traced to the fact that while 
gaseous hydrogen chloride and oxygen react 
to the extent of about 80 per cent at 400 , 
oxygen gas does not act at all on a solution of 
hydrochloric acid at room temperature. No chlorine and water, 
therefore, can be reproduced in cold water solution from the 
products of the action of these two substances, and so the 
main reaction goes on to completion. Much more is known 



Fig. 35 



Halogens with Hydrogen and Metals 143 

about the action of chlorine on water than is contained in this 
paragraph, and the subject will be taken up again in the following 
chapter. 

246. The Union of Chlorine with Metals. — Chlorine unites 
directly with many metals forming chlorides. In many cases 
the reaction takes place at once, with the production of heat and 
even in some cases of light, upon bringing the metal into chlorine 
gas. Thin pieces of copper in the form of dutch metal take fire 
when dropped into a jar of chlorine, forming copper chloride, 

Cu+Cl 2 ^CuCl 2 . 

The metal antimony (symbol Sb), in the form of powder, also 

unites with chlorine, with the production of light and heat, if 

sifted into a cylinder of the gas, antimony trichloride being 

formed, 

2Sb+ 3 Cl 2 ->2SbCl 3 . 

Chlorine also unites directly with sodium, potassium, 
magnesium, zinc, iron, aluminum, mercury, and many other 
metals to form the corresponding chlorides. 

247. The Union of Chlorine and Phosphorus. — The element 
phosphorus is a white, waxy solid which can be made from cal- 
cium phosphate, bone ash (158). We have already seen (10), 
that phosphorus burns readily in the air. In so doing it unites 
with oxygen, thus: 

4 P+50 2 ->2P 2 O s , 

forming a white, solid product, phosphorus pentoxide. Phos- 
phorus also unites directly with chlorine to form either phos- 
phorus trichloride, PC1 3 , or phosphorus pentachloride, PC1 S . 

The preparation of the trichloride may be carried out in a retort 
as shown in Fig. 36. About 20 g. of dry phosphorus are placed 
in the retort and a stream of chlorine, dried by passing it through 
a wash bottle containing concentrated sulfuric acid, is passed 
in by means of the glass tube which passes through the stopper 
of the retort. As soon as the chlorine reaches the phosphorus, 
union takes place with the formation 0} much heat and the appear- 
ance of a pale flame. The course of the reaction is readily con- 
trolled by regulating the rate of flow of the gas and by moving 



144 



Introduction to General Chemistry 



the gas inlet tube up or down in the retort. If the contents get 
too hot so that phosphorus begins to distil, the temperature can 
be lowered by raising the tube. On the other hand, if yellowish 
crystals of the pentachloride appear in the retort, the tempera- 
ture is too low and the tube should be lowered. The reaction 
occurs thus: 

2P+ 3 C1 2 ->2PC1 3 . 

Phosphorus trichloride distils over and condenses to a liquid 
in the cooled receiver. It may be purified by being distilled 




Fig. 36 

from a clean, dry retort. It is a colorless liquid which boils at 
74 . It readily unites with more chlorine, forming solid crystal- 
line pentachloride, PC1 5 : 

PC1 3 +C1 2 ->PC1 5 . 

The chlorides of phosphorus are not salts. Both compounds 
are acted upon vigorously by water, according to the following 

equations: 

PC1 3 + 3 H 2 0^H 3 P0 3 + 3 HC1 
PC1 5 + 4 H 2 0->H 3 P0 4 +5HC1. 



The products are hydrochloric acid and in the first case phos- 
phorous acid, H 3 P0 3 , and in the second case phosphoric acid, 
H 3 P0 4 (159). 



Halogens with Hydrogen and Metals 145 

248. Chlorine and Turpentine. — Turpentine is a colorless 

liquid having the formula C I0 H l6 . It reacts violently with 

chlorine, thus: 

C I0 H i6 +8C1 2 ->ioC+i6HCL 

The reaction is best shown by bringing a strip of filter paper 
which has been dipped in turpentine into a cylinder of chlorine; 
a flash of flame occurs accompanied by a dense, black smoke, due 
to the finely divided carbon formed. This reaction, as well as 
that between chlorine and water, shows the great tendency of 
chlorine to unite with hydrogen even if the hydrogen is in the form 
of a compound. 

249. Practical Uses of Chlorine. — A piece of litmus paper 
dipped into chlorine water becomes colorless. Many other 
vegetable colors are also bleached in the same way. The process 
is of great practical importance. All white cotton goods have 
been bleached by a modification of this process, which will be 
described in another chapter (351). 

In recent years a new and important use for chlorine has 
been found as a reagent for the sterilization of municipal water 
supplies. The effectiveness of chlorine is due to the fact that 
it is a powerful germicide by reason of its great chemical activity. 
The chlorine is dissolved in the water at the pumping stations 
and during the interval required for the water to flow through 
the mains it reacts with the germs present and is itself reduced 
to harmless chlorides. The water supply of the city of Chicago 
is purified in this way. 

250. The Preparation of Hydrochloric Acid. — We have 
already learned that hydrogen chloride is made by the action 
of sulfuric acid on common salt. The best laboratory method 
is that described earlier (103), the reaction taking place according 
to the following equation: 

NaCl+H 2 S0 4 ->NaHS0 4 +HCl. 

If, however, double the proportion of salt indicated by this equa- 
tion is taken and the temperature is finally raised sufficiently, 
the following reaction will take place: 

2NaCl+H 2 S0 4 ->Na 2 S0 4 +2HCl. 



146 • Introduction to General Chemistry 

By the last reaction a given quantity of sulfuric acid will produce 
double the quantity of hydrogen chloride as in the first; it is 
therefore the more economical and is the one used in the com- 
mercial production of hydrochloric acid. 

The union of hydrogen and chlorine to form hydrogen 
chloride has already been discussed (44, 243). In recent years, 
since chlorine has become available in immense quantities as 
a by-product of the manufacture of caustic soda, some hydro- 
chloric acid has been produced commercially in this way. 

The old name for hydrochloric acid was muriatic acid, and 
this is the name by which the crude acid is still commonly known 
in trade. 

251. The Physical Properties of Hydrogen Chloride. — 
Hydrogen chloride is a colorless gas, having a choking odor and 
forming a cloud of white fumes in moist air. Its density is 
somewhat greater than that of air; one liter weighs 1.642 g. 
The gas is very soluble in water; at room temperature water 
dissolves about 450 times its volume of the gas, giving a con- 
centrated solution of hydrochloric acid. Considerable heat is 
produced when the gas dissolves in water, so that the solution 
becomes decidedly warm. In general, when gases dissolve in 
water heat is produced. So-called chemically pure hydrochloric 
acid has a specific gravity of 1 . 2 and contains about 37 per cent 
of hydrogen chloride, the balance being water. 

When the 37 per cent acid is heated, hydrogen chloride 
gas is given off and the remaining solution becomes less concen- 
trated. Upon continued heating in an open vessel, the tempera- 
ture rises to no° before the liquid boils; by this time the 
concentration has decreased to 20 per cent. As the solution 
continues to boil, its concentration, 20 per cent, and boiling- 
point, no°, remain constant; the condensed vapor, the so-called 
distillate, also has a concentration of 20 per cent. 

On the other hand, if very dilute hydrochloric acid is boiled 
it loses water chiefly and becomes more concentrated; finally, 
when the concentration has reached 20 per cent the boiling 
temperature has become no°, after which both concentration 
and boiling-point remain constant. 



Halogens with Hydrogen and Metals 147 

252. The Chemical Properties of Hydrochloric Acid. — The 

most important chemical properties of hydrochloric acid have 
already been studied. These may be briefly reviewed in this 
paragraph. Hydrochloric acid is perhaps the most typical of 
all acids; it turns litmus red and its very dilute solution, say 
1 per cent, has a pleasant sour taste ; it neutralizes the hydroxides 
and oxides of metals, forming chlorides and water, for example : 

NaOH+HCl^NaCl+H 2 
CuO+ 2HCl->CuCl 2 +H 2 0. 

It acts on many metals forming chlorides and hydrogen, thus : 
Fe+2HCl^FeCl 2 +H 2 . 

The addition of hydrochloric acid to solutions of salts of 
silver (169), lead (167), and univalent mercury (182) gives 
precipitates of insoluble chlorides, thus: 

AgN0 3 +HCl->AgCl+HN0 3 . 

Oxidizing agents, such as oxygen gas at a high temperature 

and higher oxides of the metals like manganese dioxide, liberate 

chlorine : 

4 HCl+0 2 ->2Cl 2 +2H 2 (239) 

. 4 HCl+Mn0 2 ->Cl 2 +MnCl 2 +2H 2 0. (235) 

Hydrochloric acid is an almost indispensable chemical 
reagent. It is used extensively both in scientific and in technical 
work. It is manufactured in large quantities and is an impor- 
tant article of commerce. 

253. The Action of Hydrochloric Acid on Sodium Hydrogen 
Sulfate. — If concentrated hydrochloric acid is added slowly, 
with stirring, to a concentrated solution of sodium hydrogen 
sulfate, a white crystalline precipitate is formed, which, when 
filtered out, washed with a little water, and dried, is found to 
consist of pure sodium chloride. The reaction is represented 
thus: 

HCl+NaHS0 4 ->NaCl+H 2 S0 4 . 

This is seen to be just the reverse of the reaction by which 
hydrogen chloride is made from salt. It is therefore a reversible 



148 Introduction to General Chemistry 

reaction. The direction which the reaction will take depends 
upon the amount of water present and the temperature. Dry 
salt and anhydrous (water-free) sulfuric acid react practically 
completely to form hydrogen chloride and sodium hydrogen 
sulfate; while sufficiently dilute sulfuric acid and salt do not 
give off any hydrogen chloride gas. The reason is simple: the 
gas is very soluble in water, and even if it were formed it would 
remain dissolved in the water present. The fact that con- 
centrated solutions of hydrogen chloride and sodium hydrogen 
sulfate gives a precipitate of solid sodium chloride shows clearly 
that the reaction has a reversible tendency. It seems probable 
that in the presence of much water, that is, in dilute solution, 
all four of the substances are present in any solution that is made 
by bringing either pair of substances together. In such a solution 
we may say that there exists a state of equilibrium as the result 
of each pair of substances on the same side of the equation con- 
tinuously reacting to form the pair on the opposite side, thus: 

H 2 S0 4 +NaCl^NaHS0 4 +HCl. 

254. Bromine. — The element bromine (symbol Br) resembles 
chlorine more closely than does any other element. It does not 
occur free in nature. Its salts, the bromides, are frequently 
found in small amounts associated with chlorides. Sea-water 
contains a small proportion of bromides. Large quantities of 
bromides are obtained from deposits accompanying those of 
sodium nitrate in the desert regions of Chile. The brines from 
salt springs in Michigan also furnish bromides in commercial 
quantities. 

255. Sodium bromide, NaBr, potassium bromide, KBr, and 

magnesium bromide, MgBr 2 , are the commonest salts directly 

obtainable from natural salt deposits and brines. From any 

of these the element is readily set free by the action of chlorine, 

thus: 

2 KBr+Cl 2 ->2KCl+Br 2 . 

Upon passing chlorine gas into a solution of potassium bromide, 
the solution turns brown and when heated gives off reddish-brown 
vapors of bromine, which when cooled condense to liquid 



Halogens with Hydrogen and Metals 149 

bromine. Bromine is a reddish-brown liquid which has a 
density over three times that of water. It boils at 58 and readily 
volatilizes at ordinary temperatures. The vapor is, if anything, 
more irritating to mucous membranes than chlorine, and the 
liquid produces deep burns when brought into contact with the 
skin. Bromine must be handled with extreme caution. In case 
of accident wash off the bromine with water immediately; then 
consult an instructor regarding further treatment. 

Bromine dissolves in water to the extent of about 3 per cent 
to form a light-brown solution, known as bromine water. 

256. Hydrobromic Acid, HBr. — Hydrogen bromide, the 
water solution of which is known as hydrobromic acid, can be 
made by the direct union of its constituent elements: 

H 2 +Br 2 -> 2 HBr. 

The best method of making hydrogen bromine is based on the fact 
that bromine unites with phosphorus to form a tribromide or a 

pentabromide, thus, 

P+ 3 Br->PBr 3 
P+ 5 Br->PBr s . 

These compounds are entirely analogous to PC1 3 and PC1 5 (247). 
The bromides of phosphorus also resemble the chlorides in their 
reactions with water, thus : 

PBr 3 +3H 2 0->H 3 P0 3 +3HBr 
PBr 5 + 4 H 2 0->H 3 P0 4 +5HBr. 

The preparation of hydrobromic acid is carried out in the 
apparatus shown in Fig. 37. o 

Ten grams of red phosphorus, (j 

10 c.c. of water, and 20 to 25 g. 
of quartz sand are placed in a 
250 c.c. flask and 15 c.c. of bro- 
mine, contained in the dropping 
funnel, are allowed to run in 
slowly, drop by drop. The U-tube 

contains some pieces of broken 

i u • 1 • -i • . FlG - 37 

glass or brick or similar inert 

material mixed with 3 or 4 g. of red phosphorus, the object of 



150 Introduction to General Chemistry 

the glass or brick being to distribute the phosphorus so that it 
will present the maximum of surface. The hydrogen bromide 
given off is freed from accompanying bromine vapor by the 
phosphorus in the U-tube and is absorbed by water contained 
in the cylinder. The delivery tube should not dip into the 
water in the cylinder, since the gas is so soluble that there would 
be danger of water getting back into the U-tube and flask. 

257. The Properties of Hydrogen Bromide. — Hydrogen 
bromide is a colorless gas with a choking odor; it gives white 
fumes in moist air and dissolves abundantly in water to form a 
solution known as hydrobromic acid. This is a colorless liquid 
which closely resembles hydrochloric acid in its properties. It 
neutralizes bases and unites with metallic oxides to form salts 
called bromides, for example: 

NaOH+HBr^NaBr+H 2 
MgO+ 2HBr->MgBr 2 +H 2 
CuO+2HBr->CuBr 2 +H 2 
Al(OH) 3 -f 3 HBr->AlBr +3H 2 0. 

The bromides of silver, lead, and univalent mercury are almost 
insoluble in water, as are the chlorides of these same metals (252). 
All other bromides are easily soluble. The addition of hydro- 
bromic acid or any soluble bromide to a solution of a salt of 
silver, lead, or univalent mercury gives a white precipitate of 
the insoluble bromide, thus: 

Pb(N0 3 ) 2 + 2HBr->PbBr 2 + 2HNO3. 

258. The Oxidation of Hydrobromic Acid.— Hydrogen bro- 
mide and oxygen gases react when heated to form bromine and 
water, 

4 HBr+0 2 ->2Br 2 +2H 2 0. 

This reaction is analogous to that between hydrogen chloride 
and oxygen (239), but takes place far more completely, indi- 
cating that hydrogen bromide is more easily oxidized than hy- 
drogen chloride. Other oxidizing agents, such as manganese 
dioxide, readily set free bromine: 

4HBr+Mn0 2 H>MnBr 2 +Br 2 +2H 2 0. 



Halogens with Hydrogen and Metals 151 

In the technical preparation of bromine by means of this 
reaction sodium bromine is treated with dilute sulfuric acid 
and manganese dioxide. In this case' all of the available 
bromine is set free. 

2NaBr+2H 2 S0 4 H-Mn0 2 ^Na 2 S0 4 +MnS0 4 +Br 2 +2H 2 0. 

259. The Action of Chlorine on Bromides. — A solution of 

any bromide reacts with chlorine to form a chloride and free 

bromine, 

2 KBr+Cl 2 ->2KCl+Br 2 . 

Similarly, hydrobromic acid and chlorine give hydrochloric acid 
and bromine. These reactions are nearly complete, that is, 
they are not reversible to any marked extent, so that we may 
conclude that the metals and hydrogen form by preference com- 
pounds with chlorine rather than with bromine. This fact may 
also be expressed by saying that chlorine has greater affinity than 
bromine for metals and hydrogen. Using this mode of expres- 
sion, we should also say that oxygen has greater affinity than 
bromine for hydrogen, since hydrogen bromide and oxygen give 
water and free bromine. 

260. The Uses of Bromine and Its Compounds. — Potassium 
and sodium bromides are used extensively in medicine as seda- 
tives. Silver bromide is the light-sensitive substance of photo- 
graphic plates. The free element is extensively used in the 
manufacture of important coal-tar dyes. 

261. Iodine. — The element iodine (symbol I), bears almost 
the same relation to bromine that the latter bears to chlorine. 
It does not occur free in nature, but is readily prepared from 
its compounds, the iodides of sodium or potassium, which are 
obtained from two principal natural sources. 

Certain seaweeds contain small amounts of combined iodine 
which has been taken up from sea-water in which a minute 
quantity is present. The ashes left upon burning the dried 
seaweed yield by extraction with water sodium iodide, Xal, 
and potassium iodide, KI. Iodine compounds are also obtained 
as by-products in the purification of the sodium nitrate found 



152 Introduction to General Chemistry 

in Chile (104). Iodine is set free from iodides by the action of 
chlorine, thus: 

2NaI+Cl 2 ->2NaCl+I 2 . 

It is also liberated by the action of manganese dioxide and 
sulfuric acid. 

262. The Physical Properties of Iodine. — Iodine is an almost 
black, crystalline substance, having a density of nearly five. It 
melts at 114 and boils at a somewhat higher temperature, pro- 
ducing a vapor having a magnificent violet color. At a tempera- 
ture slightly below its melting-point iodine has so great a vapor 
pressure that by cautious heating it may be volatilized com- 
pletely without being melted. If the vapor is allowed to strike 
a cold surface crystals of iodine deposit directly without pre- 
liminary formation of liquid iodine. The sublimation (179) of 
iodine in this way is an important step in the purification of this 
element. 

Iodine is very slightly soluble in water, giving a faintly 
brownish solution. It dissolves abundantly in water solutions 
of potassium or sodium iodide. It dissolves easily in alcohol, 
forming a dark-brown solution called by druggists tincture of 
iodine. Iodine also dissolves easily in ether, forming a brown 
solution, and in chloroform and carbon disulfide, forming 
violet-colored solutions. 

263. Iodine and Starch. — If a dilute solution of iodine is 
added to water containing a little starch paste, made by boiling 
starch with 50 to 100 times its weight of water, a deep blue- 
colored solution results. This reaction is a characteristic and 
very delicate test for free iodine. Iodides, like KI, do not give 
this test; but by adding chlorine to a solution of an iodide the 
element is set free and can then be recognized by the starch test. 
An excess of chlorine interferes with this test. 

264. Hydrogen Iodide, HI. — Iodine and hydrogen unite 
slowly at a temperature of 400 to form hydrogen iodide, thus: 

I 2 +H 2 -> 2 HI. 

The product is a colorless gas, analogous to hydrogen chloride 
and hydrogen bromide. Like these latter gases it dissolves 
abundantly in water, and forms fumes in moist air. 



Halogens with Hydrogen and Metals 153 

Hydrogen iodide is easily made by a reaction resembling 
that used for making hydrogen bromide. Iodine forms with 
phosphorus a tri-iodide, PI 3 . This reacts with water to form 
phosphorous and hydriodic acids thus : 

PI 3 +3H 2 0->H 3 P03+3HL 

The process of making hydrogen iodide is carried out by placing 
a mixture of powdered iodine and red phosphorus in a flask and 
running in water, drop by drop from a dropping funnel, care 
being taken not to use more water than is necessary, since an 
excess of water would dissolve the gas and so prevent its escape 
from the flask. The apparatus used for making hydrogen 
bromide, Fig. 37, may be used in this case. The U-tube con- 
taining red phosphorus serves here to remove iodine vapor. 
The hydrogen iodide gas may be collected by downward dis- 
placement of air or it may be dissolved in water to form a solu- 
tion of hydriodic acid. 

265. Hydriodic acid is colorless when pure, but is brown if 
it contains free iodine, which it dissolves readily. It neutralizes 
bases and so yields salts called iodides, for example : 

HI+NaOH-^NaI+H 2 
2 HI+Ca(OH) 2 ->CaI 2 +2H 2 0. 

Hydriodic acid acts on metals similarly to hydrochloric acid, 
giving iodides and hydrogen, thus: 

2HI+Zn^ZnI 2 +H 2 . 

Hydriodic acid is much more easily oxidized than is hydro- 
bromic acid, which in turn is more easily oxidized than hydro- 
chloric acid; while all three acids are oxidized by powerful 
oxidizing agents such as manganese dioxide and lead dioxide; 
hydriodic acid, even in dilute solution, is oxidized slowly by 
atmospheric oxygen, which has no action whatever on dilute 

hydrochloric acid: 

4 HI+0 2 ->2H 2 0+2l 2 . 

The iodine which is slowly liberated according to the equation 
given above remains dissolved in the unchanged acid and gives 
it a brown color. 



154 Introduction to General Chemistry 

266. Uses of Iodine and Iodides. — Iodine is used extensively 
in certain processes of analysis and also in the preparation of 
important compounds containing the element carbon, so-called 
organic compounds. Iodine in the form of tincture of iodine, 
which is a solution of iodine in alcohol, is used externally as an 
antiseptic and also as a counterirritant in medicine. The iodides 
of potassium, sodium, and ammonium are of great importance 
for internal administration in medicine. 

267. Fluorine. — The element fluorine (symbol F), is classed 
among the halogens, although it is less closely related to the 
other three halogens, chlorine, bromine, and iodine, than these 
three are to one another. The atomic weights of these elements 
are: fluorine, 19; chlorine, 35.5; bromine, 80; iodine, 127. 
Fluorine has, therefore, the smallest atomic weight of any of 
the halogens. We might expect it to resemble chlorine more 
closely than it does bromine and iodine and, in fact, such is the 
case. It is a pale-yellow gas which is very active chemically and 
never occurs free in nature. Its most abundant natural com- 
pound is calcium fluoride or fluor-spar, CaF 2 . It also occurs as 
cryolite, sodium aluminum fluoride, 3NaF-AlF 3 . These sub- 
stances are salts of hydrofluoric acid. We might expect that 
free fluorine could be made by oxidizing hydrofluoric acid with 
manganese dioxide, thus: 

4 HF+Mn0 2 ->MnF 2 +F 2 +2H 2 0; 

but we find, in fact, that hydrofluoric acid is entirely unacted 
upon by the most powerful oxidizing agents. The free element 
was first made by Moissan, by the electrolysis of anhydrous 
liquid hydrogen fluoride, in which some potassium fluoride, KF, 
was dissolved to make it conduct electricity readily. The pro- 
ducts of the electrolysis were fluorine and hydrogen: 

2 HF->F 2 +H 2 . 

Fluorine is one of the most active of all elements. It rapidly 
attacks glass and also most metals, and it reacts at once with 
water forming hydrofluoric acid and oxygen: 

2F 2 +2H 2 0-> 4 HF+0 2 . 



Halogens with Hydrogen and Metals 155 

The preparation of fluorine is a matter of great difficulty, for 
which reason it is very seldom made. 

268. Hydrogen Fluoride, HF. — Hydrogen fluoride, a gas 
whose water solution is called hydrofluoric acid, is the most 
important compound of fluorine. It is formed by the action of 
concentrated sulfuric acid on powdered calcium fluoride : 

H 2 S0 4 + CaF 2 ->CaS0 4 + 2HF. 

It is a colorless gas with a choking odor. At temperatures of 
ioo° and higher its density shows that the gas has the formula 
HF; at room temperature the density is more than double 
that expected for a gas with the formula HF. This fact leads 
to the conclusion that the single molecules have become asso- 
ciated, probably to form double or triple molecules such as 
H 2 F 2 and H 3 F 3 . Hydrogen fluoride gas is condensed to a liquid 
merely by cooling it with ice; colorless liquid hydrogen fluoride, 
so obtained, boils at 19 . 

269. Hydrofluoric Acid and Its Salts. — A 30 per cent solu- 
tion of hydrofluoric acid is an important article of commerce. 
The acid has several practical uses. These include the etching 
and polishing of glass, the removal of sand from castings, and 
the preparation of its salts and also of hydrofluo silicic acid, 
H 2 SiF 6 . 

Hydrofluoric acid forms with bases salts called fluorides. 
The soluble fluorides are very effective preservatives, since they 
inhibit the growth of bacteria, molds, etc. But their use in 
foodstuffs is prohibited because of their interference with diges- 
tion. 

Ammonium fluoride, NH 4 F, is used as a disinfectant for 
utensils used in breweries. Sodium fluoride, NaF, is extensively 
used as a vermin exterminator for poultry. 

270. The Action of Hydrogen Fluoride on Quartz. — We must 
now digress a little from the subject in hand in order to be able 
fully to understand one of the most interesting reactions of 
hydrogen fluoride. The substance called quartz is the oxide 
of an element silicon (symbol Si) and has the formula Si0 2 . 
Common sand is more or less pure quartz. Glass, which is 



156 Introduction to General Chemistry 

made by melting together sand, sodium carbonate, and slaked 

lime, may be considered a mixture of sodium silicate, Na 2 Si0 3 , 

and calcium silicate, CaSi0 3 . Hydrofluoric acid and quartz 

react very readily to form gaseous silicon fluoride, SiF 4 , and 

water, thus: 

4 HF+Si0 2 ->SiF 4 +2H 2 0. 

This is a very characteristic reaction ; none of the other halogen 
acids have any action on quartz. 

Glass, which is almost unaffected by the other halogen acids, 
is rapidly attacked by either hydrogen fluoride gas or hydro- 
fluoric acid solution. The fluorine unites, not only with the 
silicon, as in the case of quartz, forming silicon fluoride, but also 
with the sodium and calcium forming sodium fluoride, NaF, and 
calcium fluoride, CaF 2 , the reactions being: 

Na 2 Si0 3 +6HF->SiF 4 +2NaF+ 3 H 2 
CaSi0 3 +6HF->SiF 4 +CaF 2 + 3 H 2 0. 

The result is that glass dissolves very easily in hydrofluoric 
acid, in consequence of which this acid cannot be kept in glass 
bottles. Paraffine and other waxes, which are not attacked, are 
used for bottles for this acid, while larger containing vessels are 
made of lead. 

271. Etching Glass with Hydrogen Fluoride. — The etching of 
glass may be illustrated by coating a glass plate with a thin 
layer of paraffine, and after making a design or inscription by 
means of a hard pencil which will cut through the paraffine and 
thus expose the surface of the glass, exposing the plate to the 
action of hydrogen fluoride gas. The gas is easily made by 
mixing a few grams of powdered fluor spar with concentrated 
sulfuric acid in a shallow lead dish. The latter is covered with 
the glass plate and set aside for ten or fifteen minutes. Upon 
removing the paraffine, the design will be found to have been 
etched upon the glass. 

272. Hydrofluo silicic Acid, H 2 SiF6.— Hydrogen fluoride and 
silicon tetrafluoride unite readily in the presence of water to 
form a solution of hydrofluosilicic acid: 

2 HF+SiF 4 ->H 2 SiF 6 . 



Halogens with Hydrogen and Metals 157 

The solution is a colorless, odorless liquid which does not attack 
glass appreciably. It has well-characterized acid properties: 
it reddens litmus, has a sour taste, and neutralizes bases to form 
salts. This acid is important technically. It is made, in practice, 
by the action of hydrofluoric acid solution on quartz sand : 

6HF+ Si0 2 ->H 2 SiF 6 + 2 H 2 0. 

The acid is used for the preparation of its sodium, magnesium, 
and lead salts. Sodium fluosilicate, Na 2 SiF 6 , is extensively used 
in making white enameled ware and also white, or so-called milk, 
glass. It is remarkable in being one of the very few nearly 
insoluble salts of sodium. It is obtained as a white precipitate 
when solutions of common salt and hydrofluosilicic acid are mixed. 

2NaCl+H 2 SiF 6 ->Na 2 SiF 6 + 2HCI. 

Magnesium fluosilicate, MgSiF 6 , easily soluble in water, is 
used to harden concrete. Lead fluosilicate, PbSiF 6 , also easily 
soluble in water, is made as an intermediate product in refining 
lead (Betts's process). 



CHAPTER XIII 
CHEMICAL EQUILIBRIUM 

273. Incomplete Physical Processes. — While many physical 
processes are seemingly complete, there are others which stop 
far short of completion. Thus, for example, if a small bulb of 
water is broken in a large closed bottle, evaporation of the 
water will start at once, but will apparently cease as soon as the 
pressure of the vapor reaches a value which is definite for a 
definite temperature, although much liquid may still remain 

(«2). 

If we add to some water an equal weight of common salt, 
the latter will at once start to dissolve and will continue to do so 
until the solution has, for a given temperature, a certain definite 
concentration; then, although much solid salt is still present, 
no further increase in concentration will take place (122). 

When water in a closed vessel, which it fills but partially, 
reaches its maximum vapor pressure for a given temperature, 
we believe (201) that for every molecule that passes from liquid 
to vapor there is one that passes from vapor to liquid. We say 
that there is equilibrium between liquid and vapor. We believe 
that a similar condition exists when a solid apparently stops 
dissolving in a solution (207). The apparent state of rest or 
inaction in both cases is very probably one in which two opposing 
actions exactly counteract the effects of each other. 

274. Incomplete Chemical Reactions. — Just as in the case 
of physical processes, there are also some chemical reactions 
that do not go to completion. We have already studied some 
reactions of this kind and must now consider the matter more 
fully, as it is one of great importance. 

The reaction between hydrogen chloride and oxygen at 400 
has been considered (239, 240) under the heading "Deacon's 
Process." It has been pointed out (245) that only 80 per cent 
of the hydrogen chloride is oxidized when a mixture of this gas 

158 



Chemical Equilibrium 159 

is heated with oxygen in the proportion indicated in the following 
equation : 

4HC1+0 2 ^ 2 C1 2 +2H 2 0. 

On the other hand, when a mixture of two formula weights 
each of chlorine and water is also heated to 400 , 80 per cent of 
the chlorine remains unchanged, while 20 per cent is converted 
into hydrogen chloride. It thus happens that whether we start 
with the pair of gases on the left side of the foregoing equation 
or the pair on the right, taking in each case the amounts indi- 
cated in this equation, there results a mixture of the four gases 
which has exactly the same amount of each gas present in the 
two cases. It is easy to see that the cause of each reaction being 
incomplete is found in the fact that the products of either reac- 
tion again react in the opposite direction. In the mixture of the 
four gases which finally results we say that a state of equilibrium 
exists and that the apparent stopping of further change is really 
the result of the formation of hydrogen chloride and oxygen at 
just the same rate as that at which these two gases change into 
chlorine and water. 

275. Velocity of Chemical Change. — The idea that a state 
of chemical equilibrium is the result of two opposing changes 
which take place continuously at such rates or with such velocities 
that for every molecule of a given substance formed one also 
disappears would imply that chemical changes take place gradually 
and possibly at definite speeds or velocities. 

It is well known that certain reactions, as for example the 
burning of a candle or the action of an acid on a metal, certainly 
do take place gradually. It is not so plain that if the reaction 
takes place between two perfectly mixed gases or between two 
substances completely dissolved as a uniformly mixed solution 
that time is required for the reaction to take place. Nevertheless 
it is probable that no reaction, even an explosion, however rapid 
it may be, is absolutely instantaneous. 

The speed or velocity of reaction in a uniformly mixed solu- 
tion may be beautifully and convincingly demonstrated by means 
of the following experiment: 



160 Introduction to General Chemistry 

To 800 c.c. of water contained in a flask there is added 25 c.c. 
of starch solution (made by boiling 2 g. of starch with 100 c.c. 
of water) and 15 c.c. of a 3 per cent solution of iodic acid, HI0 3 . 
The solution is then well mixed and 5 c.c. of a 3 per cent solution 
of sulfur dioxide, S0 2 , is added and the contents of the flask 
are at once thoroughly mixed by being shaken. The time of 
adding the sulfur dioxide solution is accurately noted — best with 
a stop watch. No change will be seen in the colorless solution 
for about 60 seconds, then the whole solution will suddenly turn 
deep blue. The result is startling! 

If the experiment is repeated, using the same amounts of 
water and of each of the three solutions, and if the temperature 
is also the same, it will be found that the time required for the 
change to occur is always the same. If, however, we increase 
the amount of sulfur dioxide solution added from 5 c.c. to 
10 c.c, everything else remaining the same, the time required 
for the change will be decreased to about 30 seconds. The 
increased velocity is the result of the increase in concentration of 
the sulfur dioxide. 

276. The Effect of Temperature on Reaction Velocity. — The 
effect on the velocity of increasing the temperature is easily 
shown by starting with water at 25 instead of at 20 , when it 
will be clear that at the higher temperature the velocity is decidedly 
greater. 

277. The Action of Sulfur Dioxide on Iodic Acid. — The 
chemical changes involved in the reaction just described need 
not greatly concern the student at this time, as they are of less 
importance than the main facts of reaction velocity that they 
serve here to illustrate. But as it is only natural to wonder 
what has happened in such a striking experiment, the equations 
for the reactions may now be given. In the first place, sulfur 
dioxide, S0 2 , and water form sulfurous acid, H 2 S0 3 , 

S0 2 +H 2 0->H 2 S0 3 . 

The latter reacts with the iodic acid, forming hydriodic and 
sulfuric acids, thus: 

HI0 3 +3H 2 S0 3 ->HI+3H 2 S0 4 . 



Chemical Equilibrium 161 

But hydriodic acid can also react with iodic acid to form free 
iodine and water, 

HI0 3 +5HI-> 3 I 2 +3H 2 5 

and then the iodine set free acts on the starch to produce the 
blue color. Now this reaction between iodic and hydriodic 
acid does not take place until all the sulfurous acid has dis- 
appeared. The time observed for the appearance of the blue 
color is therefore essentially that required for the complete 
oxidation of the sulfurous acid. 

278. The Kinetic Hypothesis Applied to Reaction Velocity. — 
The application of the kinetic-molecular hypothesis (chap, x) 
leads to a simple and reasonable explanation of reaction velocity. 

Let us suppose that two gases, A and B, can unite to form 

a compound AB, and let the reaction be represented by the 

equation 

A+B^AB. 

Let us also suppose that the reaction takes place rather 
slowly after the two gases have been thoroughly mixed. We 
may now consider what determines the rate at which A and B 
unite. It is obvious that union can occur only when a molecule 
of A comes in contact with a molecule of B. Such collisions 
will frequently occur by reason of the rapid motion of both 
kinds of molecules. Now as these collisions are matters of 
chance it is very easy to see that if more molecules of one or 
both kinds are brought into a given space the number of collisions 
of A molecules with B molecules will be increased. On the 
other hand, decreasing the number of one or both kinds of 
molecules will surely decrease the possible collisions of A with B 
molecules. 

Probably not every collision of an A with a B molecule will 
result in a union of the two to form AB; but if, on the average, 
a certain definite fraction of the collisions result in union, then 
we can say that the greater the number of A and B molecules 
present in a given volume, say 1 c.c, of the gas mixture, the 
greater will be the number of AB molecules formed per second. 
If we start with a mixture of equal numbers oi A and B 



162 Introduction to General Chemistry 

molecules there will be for a definite pressure and temperature a 
certain number of AB molecules formed per second. After 
a short time the number of A and B molecules will have de- 
creased appreciably, so that now fewer AB molecules will be 
formed per second, and as time goes on, owing to continual de- 
crease in the numbers of A and B molecules present, there will 
be fewer and fewer AB molecules formed per second. The 
result will be that the rate of formation of AB molecules will 
be greatest at the start and will gradually decrease, until finally, 
if the reaction is not reversible, all A and B molecules will have 
united. 

279. The Kinetic Hypothesis Applied to Chemical Equilib- 
rium. — Let us next consider, in the light of the kinetic-molecular 
hypothesis, the state of affairs if a reaction between gases is 
reversible. The case of hydrogen chloride and oxygen will serve 
as a good illustration. The equation is 

4 HC1+0 2 ^ 2 C1 2 +2H 2 0. 

This reaction takes place with moderate velocity at 400 , finally 
reaching a state of equilibrium in which all four. of the substances 
are present. 

Suppose we bring into a closed vessel at 400 a mixture of 
hydrogen chloride and oxygen in the proportion indicated by the 
equation; that is,- four molecules of the first gas to one of the 
second. The reaction will begin at a certain velocity, mole- 
cules of hydrogen chloride and oxygen disappearing by uniting 
to form molecules of chlorine and water vapor. As time goes 
on there will be fewer and fewer hydrogen chloride and oxygen 
molecules present, so that the number of each uniting per second 
and also the number of chlorine and water molecules formed per 
second will continuously decrease. On the other hand, the mole- 
cules of chlorine and water which have been formed begin to 
reunite to form hydrogen chloride and oxygen. As the total 
numbers of chlorine and water molecules present will increase 
as time goes on, so the numbers of these molecules which react 
and so disappear per second will also increase. The final result 
will be that in each second there will be just as many molecules 



Chemical Equilibrium • 163 

of chlorine and water disappearing as the numbers of each formed. 
The same sort of thing will be true for the hydrogen chloride and 
oxygen — as many molecules of each will finally be produced per 
second as the numbers that disappear. When this condition 
is reached no further change in the number of any of the four 
sorts of molecules will take place, although chemical change will 
go on continuously. The system is then in a state of equilibrium. 

We may now take up a study of a number of reversible reac- 
tions which reach a state of equilibrium. 

280. Ferric Chloride and Ammonium Sulfocyanate. — If we 
add to a very dilute solution of ferric chloride, FeCl 3 , which is 
faintly yellow in color, a dilute solution of ammonium sulfo- 
cyanate, NH 4 NCS, which is colorless, a blood-red solution results. 
This red substance is ferric sulfocyanate, Fe(NCS) 3 , which is 
formed thus : 

FeCl 3 -r-3NH 4 NCS->Fe(NCS) 3 +3NH 4 Cl. 

Let us now consider how we may discover whether this 
reaction is complete when the two substances on the left-hand 
side of the equation are mixed in the indicated proportion or 
whether a state of equilibrium results. The experiment may be 
carried out on the lecture table in the following manner: 

To 2 liters of water we add 20 c.c. of a decinormal solution of 
ferric chloride and 20 c.c. of a decinormal solution of ammonium 
sulfocyanate, which is just the amount indicated by the equation 
as required for the amount of ferric chloride present. Let us now 
divide the red solution into four equal portions, which we may 
place in four similar cylinders or beakers. Suppose we now add 
to the solution in one of the cylinders 20 c.c. more of ammonium 
sulfocyanate solution. The solution will be seen to become 
deeper red in color, which means that more red ferric sulfo- 
cyanate has been formed. Now this fact may be explained in 
either of two ways: first, that we had by mistake used, in the 
first place, less than the correct proportion of ammonium sulfo- 
cyanate indicated by the equation; or, secondly, that a state of 
equilibrium existed in the solution and that the increased con- 
centration of ammonium sulfocyanate had shifted the equilibrium 
so as to form more ferric sulfocyanate. 



164 • Introduction to General Chemistry 

We can test the truth or falsity of the first supposition very 
easily. If the original mixture contained less than the correct 
proportion of ammonium sulfocyanate, then there would be 
an excess of ferric chloride, and the addition of more of this salt 
would not increase the amount of ferric sulfocyanate and so 
increase the depth of red color. Let us add, therefore, 20 c.c. 
more ferric chloride to the solution in the second cylinder. It 
becomes deeper red! This seems to show that we are dealing 
with a condition of equilibrium as indicated by the double arrows 
of the following equation: 

FeCl 3 +3NH 4 NCS^Fe(NCS) 3 +3NH 4 Cl. 

If such is the case, then the addition of ammonium chloride to 
the solution in the third cylinder should cause a partial fading 
of the red color by reason of the partial disappearance of the red 
ferric sulfocyanate. Now this is actually what happens when the 
experiment is carried out, as can be seen by comparison with the 
color of the solution in the fourth cylinder. 

It is clear, therefore, that we have here a case of chemical 
equilibrium in which all four of the substances represented in the 
equation can coexist in the same solution. When we added more 
ammonium sulfocyanate to the solution in the first cylinder we 
increased the number of molecules of this salt and so increased 
the chances of collision of ferric chloride molecules with am- 
monium sulfocyanate molecules and this increased the number 
of ferric sulfocyanate molecules formed per second. This 
caused an increase in the total amount of the latter salt, and 
thus gave rise very quickly to a new state of equilibrium in 
which the proportion of ferric salt in the form of red sulfocyanate 
was greater than at first. 

The addition of more ferric chloride to the solution in the 
second cylinder caused a similar shift of equilibrium for anala- 
gous reasons. // is a general rule that increasing the concentration 
of either of the reacting substances on the same side of an equation 
causes a shift in equilibrium so as to form more of the substances 
on the other side of the equation. This rule is also illustrated by 
the fact that when more ammonium chloride was added to the 



Chemical Equilibrium 165 

solution in the third cylinder the color partially faded; this 
showed that some of the red ferric sulfocyanate had disappeared, 
and thus indicated that more ferric chloride and ammonium 
sulfocyanate had been formed. 

281. Hydrogen and Iodine. — We have already seen (264) 
that hydrogen unites with iodine vapor with appreciable speed 
at about 400 . The reaction is not complete, but reaches a 
state of equilibrium while there are still considerable uncom- 
bined substances present. The equation is 

H 2 +I 2 ^2ffl. 

That the reaction is reversible is easily shown by heating hydro- 
gen iodide gas, when the purple vapors of iodine appear. If the 
temperature is 370 , equilibrium is reached when one-fifth of the 
hydrogen iodide has dissociated into free iodine and free hydro- 
gen. This means that out of every 1,000 molecules of hydrogen 
iodide taken, 200 have dissociated and 800 remain when the 
state of equilibrium is reached. The equation shows that one 
molecule of hydrogen and one of iodine are formed by the disso- 
ciation of two molecules of hydrogen iodide. Therefore for 
every 200 molecules of the compound dissociated there would 
be formed 100 molecules of hydrogen and 100 of iodine. The 
equilibrium mixture resulting from every 1,000 molecules of 
hydrogen iodide taken consists, therefore, of 800 molecules of 
hydrogen iodide, 100 molecules of hydrogen, and 100 molecules 
of iodine. 

If we bring together in a closed vessel equal numbers of 
molecules of hydrogen and iodine and heat at 370 until equili- 
brium is reached we shall find that for every 500 molecules of 
hydrogen and 500 molecules of iodine taken there result 800 
molecules of hydrogen iodide, 100 molecules of hydrogen, and 
100 of iodine. In other words, just the same proportion as 
would be obtained by starting with pure hydrogen iodide gas. 

282. The Criterion of Equilibrium. — In all cases of reactions 
reaching a condition of equilibrium the resulting mixture has 
the same proportions of all substances, whether we start with 
the substances on one side of the equation or with equivalent 



1 66 Introduction to General Chemistry 

amounts of those on the other side. Therefore, if we wish to 

know whether a given reaction has reached equilibrium we bring 

together the substances which would be the products of the first 

reaction. If the resulting reaction then gives a mixture of the 

same composition as that obtained in the first case we conclude 

that both reactions have reached equilibrium. 

283. Equilibrium Constant. — In the hydrogen and iodine 

reaction 

H 2 +I 2 ^ 2 HI 

the rate of union of hydrogen and iodine, which we may 
call the speed of the reaction from left to right, will depend on 
the numbers of molecules of these two elements present in each 
c.c. It would seem probable that for a fixed number of hydrogen 
molecules per c.c. the speed of union would vary directly as the 
number of iodine molecules, and vice versa; so that this speed 
should be proportional to the product of the number of hydrogen 
molecules N x and the number of iodine molecules N 2 present in 
each c.c. of the gas mixture. That is, the speed of union, S x , is 
proportional to N x times N 2 \ or, algebraically, 

S 1 = k I XN I XN 2 , 

where k z is a constant proportionality factor. 

On the other hand, the reverse change involves the formation 
of hydrogen and iodine from hydrogen iodide, and we see by 
referring to the equation that two molecules of hydrogen iodide 
must react in order that one molecule of hydrogen and one 
molecule of iodine may be formed. This fact would make it 
seem necessary for two molecules of hydrogen iodide to collide 
in order that the change could occur. If so, increasing the num- 
ber of HT molecules in each c.c. would increase for each mole- 
cule the chances per second of collision and, in fact, doubling 
the number of molecules of this gas per c.c. would increase the 
total number of the chances per second fourfold, etc. In other 
words, the number of collisions per second of HI molecules with 

*It has become customary in chemical literature to use formulae of simple 
substances as abbreviations for the names of these substances; especially in cases 
of frequent repetition. 



Chemical Equilibrium 167 

one another will be proportional to the square of the number of 
molecules of this sort in each c.c. The details of the method of 
arriving at this conclusion need not be considered at present. 
If we call the speed of change of hydrogen iodide into hydrogen 
and iodine 5 2 and call the number of HI molecules in 1 c.c. N 3J 
then it is plain that this speed is proportional to N 3 2 , or 

S 2 = k 2 N 3 2 , 

where k 2 is a constant proportionality factor. 

Let us now think of the state of affairs when equilibrium has 
resulted. The speed of formation of hydrogen iodide which 
is equal to the speed of union of hydrogen and iodine, S lt is now 
just equal to the speed of dissociation, 5 2 , of the hydrogen iodide. 
This must be the case, as otherwise further changes in the pro- 
portions of the three substances would still be taking place and 
the mixture would not be in equilibrium. For the state of 
equilibrium, therefore, we may write 

and hence 

hXN.XN^hXN, 2 
or 

N 3 2 = k x 
NtXN 2 k 2 ' 

Now kt and k 2 are both constant quantities for the reaction 
under consideration if the temperature is fixed, and therefore 
their quotient is a constant, so that we may write 



a constant. Therefore, 



fV2 . 



N 2 

iV3 =K. 



NtXN. 



This algebraic equation means that for the condition of equilibrium 
at a fixed temperature the square of the number of molecules per c.c. 
of HI divided by the product of the numbers of molecules of H 2 and 
1 2 is a fixed or constant quantity. This matter can perhaps be 
made a little plainer by use of a numerical example. We have 



1 68 Introduction to General Chemistry 

seen that at 370 the equilibrium mixture which results from 
1,000 original HI molecules consists of 800 molecules of HI, 100 
of H 2 , and 100 of I 2 . In each c.c. of such an equilibrium mixture 
the total number of molecules will be very great; but, of course, 
the numbers of each kind will be in the same proportion as for a 
total of 1,000 molecules, and therefore 

100X100 

If we start with unequal instead of equal numbers of molecules 
of hydrogen and iodine we can calculate by means of the equa- 
tion 

N 2 

mxwr 64 

what the state of equilibrium will be. For example, suppose we 
start with a mixture of hydrogen and iodine containing four 
times as much hydrogen as would theoretically be necessary 
for the iodine taken; that is, four molecules of hydrogen for one 
of iodine. Calculation shows that if we start with 800 mole- 
cules of hydrogen and 200 molecules of iodine, when equilibrium 
is reached, out of a total of 1,000 molecules 392 will be hydrogen 
iodide, 604 will be free hydrogen, and 4 will be free iodine. 

284. Ammonia and Water. — Several reactions already studied 
reach a condition of equilibrium; three of the most familiar of 
these may now be considered as additional examples of the subject 
under discussion. Ammonia gas, NH 3 , dissolves abundantly in 
water, giving a solution which turns litmus blue and forms salts 
with acids. The solution contains ammonium hydroxide, 
formed by the union of ammonia with water (91) : 

NH 3 +H 2 0->NH 4 OH. 

The solution smells strongly of ammonia and if it is boiled a 
short time all of the gas is given off. This shows that ammonium 
hydroxide easily dissociates into its constituents. It seems 
highly probable that in the water solution a condition of equilib- 
rium exists, as indicated in the equation 

NH 3 -f-H 2 O^NH 4 OH, 



Chemical Equilibrium 



169 



both free ammonia and ammonium hydroxide being present. 
Heating such a solution renders the free ammonia less soluble, 
and as this partially escapes, the rate of formation of ammonium 
hydroxide falls farther and farther behind the rate of dissociation 
of this compound until finally all of the latter has disappeared. 

285. Carbon Dioxide and Water.— A water solution of carbon 

dioxide, C0 2 , contains carbonic acid, H 2 C0 3 . But such a solution 

easily gives off carbon dioxide, especially if warmed; which leads 

us to conclude that the reaction is a reversible one, and that in 

the solution there is a state of equilibrium as represented by 

the equation 

C0 2 +H 2 O^H 2 C0 3 . (152) 

286. Sulfur Dioxide and Water. — Sulfur dioxide, S0 2 , which 
is formed when sulfur burns, is a colorless gas with a suffocating 
odor: 

S+0 2 ->S0 2 . 

It is easily soluble in water, giving a solution which smells 
strongly of the gas and has acid properties. The solution con- 
tains a compound, sulfurous acid, H 2 S0 3 . This solution gives off 
all of its sulfur dioxide when boiled, and we conclude, therefore, 
that the acid easily decomposes into its constituents, water and 
sulfur dioxide, and that in the solution we have a state of equilib- 
brium, as- represented by the equation 

S0 2 +H 2 O^H 2 S0 3 . 

287. The Effect of Pressure on a System in Equilibrium. — 

Suppose we have, say, 1 liter of water saturated 
with a gas, say oxygen, at a fixed temperature 
and at one-atmosphere pressure. To say that 
the water is saturated with the gas means that 
a condition of equilibrium exists between solu- 
tion and gas. Let us suppose the solution and 
gas are contained in a cylinder fitted with a gas- 
tight piston (Fig. 38) and that the volume of the 
undissolved oxygen gas above the solution is 1 
liter. If now we double the pressure on the gas 
more of the gas passes into solution, finally producing a new 



Fig. 38 



170 Introduction to General Chemistry 

state of equilibrium. By reason of the fact that part of the 
gas dissolved when the pressure was doubled the volume of the 
remaining gas will not be half a liter, as we should expect if no 
additional quantity of oxygen dissolved in the water present, 
but appreciably less than half a liter. The effect, therefore, of 
increasing the pressure on the system is to cause its volume to 
become smaller than would be the case if no shift of equilibrium 
had occurred. This is the way in which an increase of pressure 
always affects a system in equilibrium : the state of equilibrium 
shifts in such a way as to cause a greater decrease in volume than 
would be the case if no change in the state of equilibrium occurred. 
Let us consider another case. We may inquire how the 
equilibrium represented by the equation 

4 HC1+0 2 ^2C1 2 + 2 H 2 

would be affected by increase of pressure. We see by reference 
to the equation that four volumes of HC1 and one of 2 give two 
volumes of Cl 2 and two of H 2 0; that is, that when the reaction 
takes place from left to right there is a decrease in volume from 
5 to 4. We should expect, therefore, that by increasing the 
pressure the equilibrium would shift somewhat from left to 
right; that is, that more chlorine and water would be formed 
at the expense of the hydrogen chloride and oxygen; and this 
is exactly what actually happens. 

The effect of increase of pressure- on any system in equilibrium 
is, in all cases, to shift the equilibrium so as to favor the formation 
of substances occupying a smaller volume. In case no change of 
volume accompanies a chemical reaction, then the state of 
equilibrium is not affected by change of pressure. The reaction 

H 2 +I 2 ^ 2 HI 

is an example of this sort. Here one volume of hydrogen and 
one volume of iodine vapor react to form two volumes of hydro- 
gen iodide, so that no change of volume occurs when the reaction 
takes place. It has been found by careful investigation that 
the equilibrium proportion of the three substances is not changed 



Chemical Equilibrium • 171 

by altering the pressure, as long as the. temperature remains 
constant. 

288. Effect of Temperature on a System in Equilibrium. — We 
have already learned (112) that the vapor pressure of water 
increases with increase of temperature. We know also that a 
large amount of heat is absorbed when water is evaporated; at 
ioo° it requires 540 calories to change one gram of water into 
steam. This is the so-called latent heat of vaporization. If 
we have, in a closed vessel, water in equilibrium with its vapor, 
and then apply heat, two effects are produced : the temperature 
is raised and the vapor pressure is increased. The increase in 
pressure is caused by the evaporation of some water, and this 
evaporation absorbs some of the heat which has been applied. 
This is a typical case, for we always find that when we apply heat 
to any system in equilibrium that the state of equilibrium shifts 
in such a way that heat is absorbed in the change. As heat is 
absorbed when water evaporates, heating causes increased vapor 
pressure. 

The effect of temperature on the solubility of substances 
has already been studied (134). We have learned that heat is 
either absorbed or produced when a substance dissolves; this is 
the so-called heat of solution. Substances which dissolve with 
absorption of heat become more soluble with rise of tempera- 
ture, while those which dissolve with evolution of heat, like 
anhydrous sodium sulfate, Na 2 S0 4 , decrease in solubility as the 
temperature is raised (134, Fig. 27). If a substance like the 
last named dissolves with evolution of heat, its crystallization 
out of a solution is accompanied by absorption of heat. In 
every case raising the temperature causes that change of solubility 
to occur which involves an absorption of heat. 

The state of chemical equilibrium is shifted in all cases by a 
change of temperature. Now we find that every chemical reac- 
tion either gives out or absorbs heat. When substances burn, the 
heat given out is very great. In many other reactions the heat 
produced is considerable, while in still others an absorption of 
heat occurs. If *a reaction is reversible (all reactions that reach 
a state of equilibrium are, of course, of this class) and produces 



172 Introduction to General Chemistry 

» 
heat when it goes in one direction, it absorbs an equal amount of 
heat for the same quantity of materials transformed when it goes 
in the opposite direction. 
If the reaction 

H 2 +I 2 ^2HI (264) 

has reached a state of equilibrium at 370 , out of every 1,000 
molecules present 800 will be HI, 100 H 2 , and 100 I 2 . If the 
temperature is then raised to 440 and held constant until a 
new state of equilibrium is reached, the gas mixture will consist 
of 780 molecules of HI, no of H 2 , and no of I 2 . Part of the 
HI has changed to H 2 and I 2 , and the equilibrium may be said 
to have shifted from right to leftr At temperatures between 
370 and 440 the change of HI into H 2 and I 2 takes place with 
absorption of heat. We see, then, that raising the temperature 
causes the equilibrium to shift in the direction that involves an 
absorption of heat. Now this is a perfectly general law for 
chemical changes, just as it is also for physical changes like the 
vaporizing of a liquid and dissolving of a solid. 

289. The Effect of Removing One Product of a Reaction. — 
The reaction represented by the equation 

NaCl+H 2 S0 4 ^NaHS0 4 +HCl (103, 253) 

has already been studied rather fully. We may summarize 
the facts briefly, as follows: The action of concentrated sulfuric 
acid on dry salt gives sodium hydrogen sulfate, NaHS0 4 , and 
hydrogen chloride gas, the reaction going nearly to completion 
in the direction of the lower arrow in the equation given above 
if the mixture is warmed. On the other hand, if a cold saturated 
solution of sodium hydrogen sulfate is mixed with concentrated 
hydrochloric acid — that is, a saturated solution of hydrogen 
chloride in water — an abundant precipitate of solid salt, NaCl, is 
formed. This reaction is, we see, just the reverse of the other. 
If now we mix a dilute solution of salt with dilute sulfuric acid, 
we see no visible change. We also see no change upon mixing 
a dilute solution of sodium hydrogen sulfate with dilute hydro- 
chloric acid. 



Chemical Equilibrium 173 

We are now in position to explain all the facts of the foregoing 
paragraph from the standpoint of chemical equilibrium. If 
we bring together dilute solutions of either pair of substances 
in the reaction 

NaCl+H 2 S0 4 ^NaHS0 4 +HCl, 

the resulting solution probably contains all jour substances, side 
by side, in a state of equilibrium. But we cannot notice any 
effect of the mixing, because in the presence of much water all 
four are held completely in solution, since all four are more or 
less readily soluble in water. If, however, but little water is 
present, the least soluble of the four substances, common salt, 
may partially separate. This is the case when a concentrated 
solution of NaHS0 4 is mixed with concentrated HC1. The 
reason is a simple one: the substances taken react partially to 
form NaCl and H 2 S0 4 in the sense of the upper arrow; but the 
amount of NaCl so formed is more than the water present can 
hold in solution; so the excess NaCl separates out in the solid 
form. This separation of NaCl continues until the four sub- 
stances in the solution have reached amounts which can and do 
exist in equilibrium with one another. Removing the solid 
NaCl which has separated, or adding more solid salt, will in no 
way alter the amounts of any of the four substances contained in 
the solution. 

When concentrated H 2 S0 4 is mixed with dry NaCl, NaHS0 4 
and HC1 begin to be formed. Now HC1 is but slightly soluble 
in concentrated H 2 S0 4 and, being a gas, it at once escapes from 
the mixture. Warming the mixture also promotes the escape 
of the HO, since the higher the temperature the smaller the 
solubility of the gas in the concentrated H 2 S0 4 . The escape of 
the HC1 gas also has another fundamental effect on the reaction. 
In order that any reaction may reach a state of equilibrium it 
must be reversible; but this reaction cannot go in the reverse 
direction if the HC1 escapes from the reacting mixture as fast 
as it is formed. The result is that if no water is present, con- 
centrated H 2 S0 4 and dry NaCl react practically completely, 
giving solid NaHS0 4 and HC1 gas. 



174 Introduction to General Chemistry 

290. The Action of Steam on Iron and the Reverse Action. — 

When steam is passed over heated iron (29, Fig. 16) hydrogen 
and an oxide of iron are formed. On the other hand, if hydrogen 
is passed over the heated oxide, Fe 3 4 , the products are iron and 
water. The equation for these two reactions, of which one is 
the reverse of the other, is 

3Fe+ 4 H 2 O^Fe 3 4 +4H 2 . 

If we bring together either pair of substances in a closed 
vessel and heat them for some time, a state of equilibrium is 
reached in which all four substances are present. The iron and 
iron oxide are solids, while the water, as steam, is a gas. For 
the condition of equilibrium the relative amounts of steam and 
hydrogen are always the same for a given temperature, no matter 
what proportions of either pair of substances have been used. 
This is the state of affairs if the reaction occurs in a closed vessel. 
But the results are entirely different if the reactions take place 
in a tube holding the solids, through which either steam in the 
one case or hydrogen in the other is passed. If steam is passed 
through a tube containing iron, then the hydrogen which is 
formed is carried along with the excess of steam and has no 
chance to act on the iron oxide which has been formed. There 
is therefore no chance for iron to be formed again, once it has 
been changed to iron oxide. As long as unchanged iron remains 
and the current of steam is continued, the reaction from left to 
right continues. The inevitable result is the complete change 
of the iron to the oxide. On the other hand, if a current of 
hydrogen is passed over heated iron oxide contained in a tube, 
the substances react in the direction from right to left of the 
equation. The steam which is formed passes along with the 
excess of hydrogen, and once having left the tube cannot pos- 
sibly act on the iron to convert it back into oxide, so that this 
change also continues as long as the stream of hydrogen is kept 
up and comes to an end only when all of the iron oxide has been 
reduced to metallic iron. 

291. Conclusions. — We see, therefore, that a chemical 
reaction like the one just discussed may reach a state of equilib- 



Chemical Equilibrium 175 

rium, if the reverse reaction tends to take place noticeably, and 
if none of the substances involved escape from the vessel in 
which the change takes place ; or it may go to completion in one 
direction or the other if one of the products of either reaction is 
allowed to escape from the scene of action. 

Whether a given reaction reaches a state of equilibrium or 
goes to completion in one direction or the other often depends 
upon the conditions. In the preparation or manufacture of 
chemical substances it is usually very important to cause equi- 
librium reactions to take place as completely as possible in 
order to obtain the maximum possible yields of the desired 
products. 



CHAPTER XIV 
HYDROGEN AND OXYGEN 

292. Hydrogen. — -Hydrogen was first recognized in 1766 as 
a distinct substance by Cavendish, a celebrated English chemist, 
who called the gas inflammable air and prepared it by the action 
of acids on metals. It was not until ten years after Cavendish's 
discovery that Lavoisier explained the role played by oxygen 
in combustion and stated the law of the indestructibility of 
matter (21) and thus laid the foundation for the doctrine of the 
elements in its present form. For this reason the classification 
of hydrogen as an element was not possible at the time of its dis- 
covery. In 1 781 Cavendish showed that nothing but water is 
formed when hydrogen burns and thus proved that water is a 
compound of hydrogen and oxygen. The name hydrogen means 
water-former. 

The element occurs in but minute amounts in the free form 
in nature. Water is its most abundant compound; but it is also 
a constituent of all dry animal and vegetable tissues, forming 
therein principally compounds with carbon, oxygen, and nitrogen. 
Petroleum and natural gas are compounds of hydrogen with 
carbon; coal also contains considerable combined hydrogen. 

293. Preparation of Hydrogen. — We have already learned 
several methods by which free hydrogen can be obtained. These 
may now be briefly reviewed. Hydrogen is formed: 

1. By the electrolysis of water (27), 

2 H 2 0->2H 2 +0 2 ; 

2. By the action of water on some metals, as by (a) the burn- 
ing of magnesium wire in steam (28, Fig. 15), 

Mg+H 2 0->MgO+H 2 , 

(b) the passage of steam over heated iron turnings (29, Fig. 16), 

3 Fe+4H 2 O^Fe 3 4 +4H 2 , 

176 



Hydrogen and Oxygen 



177 



(c) the action of sodium or potassium on water (40, 86, Table VI, 

106), 

2 Na+ 2 H 2 CM>2NaOH+H 2 
2K+2H 2 0->2KOH+H 2 ; 



3. By the action of hydrochloric or sulfuric acid on zinc, 
magnesium, iron, or aluminum, as well as on several other metals, 



Zn+ 2 HCl->ZnCl 2 +H 2 
Fe+H 2 S0 4 ->FeS0 4 +H 2 . 



(149) 
(i73) 



294. Making Hydrogen in the Laboratory. — The best labora- 
tory method of making hydrogen consists in treating zinc with 
hydrochloric acid in some form of 
specially constructed gas generator. 
The Kipp apparatus, Fig. 39, is the 
form most extensively used. The 
solution used is made from equal 
volumes of concentrated hydrochloric 
acid and water. The action of this 
generator is very simple in principle. 
Upon opening the stopcock gas 
escapes and allows the acid to rise 
into the middle compartment, where 
it acts upon the zinc and so produces 
a steady flow of hydrogen. When 
the cock is closed the gas formed 
forces the acid downward and causes it to flow from the 
lower into the upper compartment. As soon as the acid is out 
of contact with the zinc all action stops, and no more gas is 
produced until the cock is again opened. The Kipp apparatus 
has one unfortunate defect: since there is but little circulation 
of the solution the acid in contact with the zinc is soon 
exhausted, causing the action to stop while there is still a 
large supply of almost unchanged acid in other parts of the 
apparatus. To start the action again it is necessary to empty 
all the solution and refill with fresh acid; much 'acid is thus 
wasted. 




Fig. 39 



1 7 8 



Introduction to General Chemistry 



The McCoy apparatus, shown in Fig. 40, has several ad- 
vantages over the Kipp apparatus. The lowest compartment 
is filled as full as possible with granulated or stick zinc, on which 
hydrochloric acid drops at just the rate required to keep up the 
stream of hydrogen that is being drawn from the apparatus. 
When the stopcock is closed the gas which is formed from the 
small excess of acid in the zinc compartment forces the acid from 
the middle to the upper compartment and thus stops the further 
flow of acid upon the zinc. This appa- 
ratus is also conveniently used for 
generating other gases. 

295. The Electrolysis of Water. — 
Compared with metals, pure water is a 
very poor conductor of electricity. The 
addition of a little sulfuric acid increases 
the electrical conductivity of water 
enormously. The sulfuric acid so added 
is not permanently destroyed in the 
course of the electrolysis (27), so that 
very little will serve to promote the 
electrolysis of a large amount of water. 
The exact way in which the acid behaves will be discussed 
later. Ordinarily, poles or electrodes of the elementary metal 
platinum are employed, since most other metals would be 
attacked chemically. The electrode at which the hydrogen is 
liberated is called the negative electrode, or cathode ; the other, 
at which the oxygen appears, is the positive electrode, or 
anode. 

One of the important technical methods of making hydrogen, 
which yields at the same time oxygen, consists in the electrolysis 
of water in which sodium hydroxide is dissolved to make it a 
good conductor. Here the cathode is of iron and the anode of 
carbon. Hydrogen is also obtained in commercial quantities 
as a by-product in the manufacture of caustic soda by the elec- 
trolysis of a solution of common salt. 

Hydrogen is often made for use in balloons by the action of 
dilute sulfuric acid on scrap iron. 




Fig. 40 



Hydrogen and Oxygen 179 

296. The Physical Properties of Hydrogen.— We have already 
learned that hydrogen is colorless ; when perfectly pure it is also 
odorless and tasteless. One liter of the gas at o° and 76 cm. pres- 
sure weighs 0.0899 g.; an d 22.4 liters, 2 g. approximately. It 
is the lightest of all gases. It can be liquefied, giVing a colorless 
liquid which boils at — 253 , or only 20 above absolute zero. At a 
somewhat lower temperature the liquid freezes to a colorless 
solid. 

Hydrogen is but slightly soluble in water: 100 c.c. of water 
dissolves about 2 c.c. of the gas at room temperature. 

The speed of diffusion of hydrogen is greater than that of 
any other gas (191). 

297. The Chemical Properties of Hydrogen. — The most 
important chemical properties of hydrogen have already been 
studied, but may now be briefly reviewed. Hydrogen burns with 
an almost non-luminous flame, which is, however, very much 
hotter than that obtained from ordinary fuel or illuminating gas. 
Water is the product of the reaction. Hydrogen reacts readily 
with hot copper oxide, forming water and copper: 

H 2 -f-CuO->H 2 0+Cu. (33) 

Hydrogen also acts on other metallic oxides at a red heat, for 

example : 

4 H 2 +Fe 3 4 -> 4 H 2 0+3Fe. (290) 

Hydrogen and chlorine, if mixed in equal volumes and ignited, 

or exposed to a bright light, unite with explosive violence, forming 

hydrogen chloride, 

H 2 +C1 2 ->2HC1. (243) 

A jet of burning hydrogen lowered into a jar of chlorine con- 
tinues to burn by reason of the union of the two elements (244) . 
Hydrogen unites with bromine to form hydrogen bromide 
(256) and with iodine to form hydrogen iodide (264). 

298. The Union of Hydrogen and Nitrogen. — A mixture of 
hydrogen and nitrogen does not react at all under ordinarv 
conditions. If electric sparks are passed through the mixed 



i8o 



Introduction to General Chemistry 



gases contained in a eudiometer, Fig. 41, a small amount of 
ammonia is formed : 

3 H 2 +N 2 -> 2 NH 3 . 

The reaction soon reaches a state of equilibrium, because under 
the same conditions ammonia is very largely decomposed into 
its elements. As the result of the reverse reaction, a state of 
equilibrium is reached when less than 1 per cent of the ele- 
mentary gases has been converted into ammonia. If the 
ammonia is absorbed in some suitable way, 
as by union with sulfuric acid, as fast as it is 
formed, then, with continued sparking, the 
formation of ammonia goes on until all 
the hydrogen and nitrogen have united. The 
practical method of making ammonia by this 
reaction will be discussed in chapter xxi. 

299. Heat of Reaction and Flame 
Temperature. — When 1 g. of hydrogen burns, 
about 34,000 calories of heat are produced; 
this is sufficient to heat 340 c.c. of water 
from o° to ioo°. By reason of the great 
amount of heat produced, the flame of hydrogen burning in air 
has a very high temperature. When hydrogen burns in pure 
oxygen instead of in air, the flame is much hotter, but not 
because a greater amount of heat is produced by the burning 
of a given amount of hydrogen, since the quantity of heat is 
the same in the two cases. When hydrogen burns in air, the 
nitrogen, which forms four-fifths by volume of the air, is 
heated to the flame temperature along with the steam formed. 
But in pure oxygen no nitrogen is present, and so the tempera- 
ture reached by the flame is much higher, as there is far less 
material to be heated. 

300. The Oxyhydrogen Blowpipe. — Fig. 42 shows an oxyhy- 
drogen blowpipe. The two gases mix in the proper proportions 
before issuing from the jet. The temperature of the flame is high 
enough to melt platinum, which cannot be melted in a Bunsen 
flame supplied with fuel or illuminating gas. 




Fig. 41 



Hydrogen and Oxygen 181 

301. The Limelight. — A very bright light is produced when 
an oxyhydrogen flame strikes a stick of quicklime, by reason of 
the bright white heat to which the lime is raised. This is the 
so-called limelight, which was very extensively used before the 
electric arc light was perfected and which is still frequently 
used in rural communities. In place of lime other difficultly 
fusible white oxides may be employed. For this purpose 
thorium oxide containing 1 per cent of cerium oxide is much 
superior to lime. 

302. Ignition Temperature. — A mixture of hydrogen and 
oxygen in their combining proportion remains unchanged for any 




Fig. 42 

length of time at room temperature, but if brought in contact 
with a flame or electric spark it explodes violently. The explo- 
sion is due to the great increase in volume of the reaction product, 
steam, caused by the almost instantaneous reaction, with its 
attendant heat production. But we may inquire why a reaction 
which does not take place at room temperature can become explo- 
sive. Investigation shows that a mixture of the two gases reacts 
perceptibly at 450 , and that the formation of water goes on 
faster the higher the temperature, but that the mixture does not 
become explosive until the temperature reaches about 6oo°. It 
is easy to see why explosion finally occurs when the temperature 
is raised. While the reaction is taking place slowly, heat is being 
produced by the chemical change; below 6oo° the rate of change 
is so slow that heat is lost by the gas mixture faster than it is 
produced. Above 6oo° the reaction goes faster, so that heat is 
produced more rapidly than it is lost, and this causes the gas 
mixture to grow hotter; and the hotter it gets the faster the 
reaction goes, until soon it proceeds with enormous rapidity, and 
this constitutes an explosion. For any combustible substance 
there is some temperature to which it must be heated before its 



182 Introduction to General Chemistry 

rate of production of heat by union with oxygen exceeds its rate 
of loss of heat; if heated to this temperature the substance takes 
fire and continues to burn. This point is called the ignition 
temperature. 

Hydrogen, issuing from a jet, burns quietly in air when 
ignited. This is because the actual union with oxygen can occur 
only as fast as the two gases can reach one another by diffusion 
(191), one from the jet, the other from the surrounding air. 
The flame is the reacting gas mixture, which is raised to incan- 
descence by the great heat produced by the union. 

303. Platinum as a Catalytic Agent. — The element platinum 
can be deposited on asbestos as a thin, spongy coating by dipping 
a bit of fibrous asbestos in a solution of platinum chloride, 
PtCl 4 , drying the material and holding it in a Bunsen flame for 
a minute. The salt decomposes into its elements thus: 

PtCl 4 ^Pt+2Cl 2 . 

If .a jet of cold hydrogen gas is directed on the cold platinized 
asbestos, the latter gets red-hot and sets fire to the hydrogen. 
Spongy platinum absorbs gases to a marked extent. Heat is 
produced in this way, and this ignites the intimate mixture of 
hydrogen and oxygen condensed on the surface of the metal. 
The platinum itself is entirely unchanged and will continue 
active in this way indefinitely. A substance which initiates or 
promotes a chemical reaction without itself being changed is 
called a catalytic agent (239). 

304. The Use of Hydrogen in Balloons. — The lifting power 
of a balloon filled with hydrogen can easily be calculated. Since 
1 liter of air under standard conditions weighs 1 . 29 g. and 
1 liter of hydrogen weighs 0.09 g., the difference, 1.2 g., repre- 
sents the lifting power per liter of capacity of a balloon. At 
higher temperature and lower pressure the lifting power is 
smaller. If a Zeppelin has a capacity of 5,000,000 liters its 
lifting power will be about 6,000 kilos, or more than 13,000 
pounds. 

305. Oxygen. — Oxygen in the form of compounds makes up 
about one-half by weight of the matter forming the crust of the 



Hydrogen and Oxygen 183 

earth. It also constitutes 89 per cent by weight of water and 
21 per cent by volume of the air. The oxygen of the air is not 
chemically combined but is only mixed with nitrogen and small 
amounts of other gases present. Oxygen was first prepared by 
Priestley, in England (1774), by heating mercuric oxide, 

2 HgO->2Hg+0 2 . 

At practically the same time Scheele, in Sweden, made oxygen 
by this method and also by heating potassium nitrate, 

2KN0 3 ->2-KN0 2 +0 2 . 

The salt KN0 2 is called potassium nitrite. The name oxygen, 
which means acid-former, was given to the gas by Lavoisier, who 
believed that it was a necessary constituent of all acids. At 
that time hydrochloric acid, which was called muriatic or marine 
acid, was thought to contain oxygen. We now know many 
other acids which do not contain oxygen. 

306. The Preparation of Oxygen. — We have already learned 
several ways by which oxygen may be made. The heating of 
mercuric oxide and the electrolysis of water (14, 295) have already 
been fully studied. We have also seen (245) that oxygen is 
formed when chlorine water is exposed to sunlight: 

2 Cl 2 +2H 2 0->0 2 +4HCl. 

The heating of certain salts which are rich in oxygen is also 
a simple way of making the gas. The behavior of potassium 
nitrate, KN0 3 , is given in the preceding paragraph. Potassium 
chlorate, KC10 3 , is easily decomposed by heat according to the 
following equation: 

2KC10 3 ->2KCH- 3 2 . 

This last reaction is the one usually employed in making small 
amounts of oxygen in the laboratory. It can be carried out in 
a test tube, a small flask, or a retort. The crystals of potassium 
chlorate first melt, and at a little higher temperature the liquid 
seems to boil, by reason of the oxygen given off. 



184 Introduction to General Chemistry 

The change of the chlorate, KC10 3 , into chloride, KC1, does 
not take place completely in one step. The first stage of the 
reaction is probably represented by the equation, 

ioKC10 3 ^ 4 KCl+6KC10 4 +30 2 . 

The salt KC10 4 , called potassium perchlorate, can also be decom- 
posed by heat, thus : 

KC10 4 ^KCl+ 2 2 . 

This last reaction requires a higher temperature than the first. 
If the heating of the chlorate is stopped when about one-fifth of 
its total oxygen has been given off, KC10 4 will be found in the 
residue. 

In making oxygen from potassium chlorate two precautions 
should be observed : first, the material must be entirely free from 
bits of wood, paper, etc., which are easily combustible; and, sec- 
ondly, the heating must be gentle, as otherwise the decomposition 
may occur explosively. 

When powdered potassium chlorate is mixed with about half 
its weight of manganese dioxide, Mn0 2 , it will give off its 
oxygen rapidly at a temperature far below that at which the 
pure chlorate starts to decompose. Since manganese dioxide 
alone does not give off any of its oxygen until a rather high 
temperature is reached, and is not changed itself in promoting 
the decomposition of the potassium chlorate, we must consider 
that the former substance acts only as a catalytic agent in pro- 
moting the decomposition of potassium chlorate. 

307. Oxygen from Sodium Peroxide. — Sodium peroxide, 
Na 2 2 , is a solid made by burning metallic sodium, 

2Na+0 2 ^Na 2 2 . 

The trade name of the material is oxone; it is supplied in the 
form of lumps or sticks. Water acts on it as follows: 

2Na 2 2 + 2H 2 0->4NaOH+0 2 . 

By dropping water on lumps of oxone contained in a suitable 
apparatus (236, Fig. 30) a steady stream of oxygen is obtained. 



Hydrogen and Oxygen 185 

The method is rather expensive, but it is very convenient, since 
the action stops when the supply of water is turned off and can 
be started again at will. 

308. Oxygen from Other Oxides. — Lead dioxide, Pb0 2 , when 

strongly heated gives oxygen and lead monoxide, or litharge, 

PbO: 

2Pb0 2 ->2PbO+0 2 . 

Manganese dioxide is also decomposed at a high temperature, 

thus: 

3Mn0 2 ->Mn 3 4 +0 2 . 

309. Technical Methods of Making Oxygen. — The elec- 
trolysis of water is an important technical method of making 
oxygen. It also yields hydrogen and has already been described. 
By far the larger part of the oxygen of commerce is made from 
liquid air. This substance is a mixture of liquid oxygen and 
liquid nitrogen. The latter boils about n° lower than the 
former, whose boiling-point is — 183 , and therefore distils off 
first when liquid air is allowed to evaporate, leaving nearly pure 
oxygen. This is stored under pressure in steel tanks and brought 
on the market. 

310. Brin's Process. — Brin's process, formerly used tech- 
nically, is a method of obtaining oxygen from the air by means of 
barium oxide, BaO. This oxide unites with more oxygen at a 
red heat, forming barium peroxide : 

2BaO+0 2 ±5 2Ba0 2 . 

This is a reversible reaction, which at a constant temperature 
will go in one direction or the other with change of pressure. 
In practice, air is pumped under pressure into a vessel contain- 
ing BaO at 700 , and when all the oxide has been changed into 
Ba0 2 the nitrogen present is allowed to escape. By reducing 
the pressure with a vacuum pump the Ba0 2 is caused to decom- 
pose completely, yielding nearly pure oxygen. 

311. Oxygen from Plants. — Growing plants absorb carbon 
dioxide from the air. They also take up water through their 
roots. In some manner, not fully understood, carbon dioxide 
and water react under the influence of sunlight to form such 



1 86 Introduction to General Chemistry 

principal plant constituents as starch, cellulose, and sugar, 
together with oxygen, which is given off to the air. The per- 
centage of oxygen in the air would soon decrease if it were not 
maintained by growing plants. 

312. The Physical Properties of Oxygen. — It is, of course, 
obvious that oxygen is colorless, odorless, and tasteless. One 
liter weighs 1.429 g. and 22.4 liters about 32 g., corresponding 
to the formula 2 . Liquid oxygen is pale blue in color; it 
boils at — 183 . At o°, 100 c.c. of water dissolves about 5 c.c. 
of oxygen; at 20 , about 3 c.c. (125). 

313. The Chemical Properties of Oxygen. — We have already 
learned that combustion was first explained by Lavoisier in 
1774 as due to union of the burning substance with the oxygen 
of the air (13-15). All the elements so far studied, except 
fluorine, form oxides. This does not mean that all these ele- 
ments burn, since some oxides, like those of chlorine and silver, 
can only be made indirectly (172). The oxides of metallic 
elements, by union with water, form hydroxides which are 
bases, for example: 

CaO+H 2 0->Ca(OH) 2 . 

The oxides of non-metallic elements, including carbon, sulfur, 
nitrogen, phosphorus, and the halogens (except flourine), give, 
with water, acids. The following equations will serve as illus- 
trations of such reactions, some of which have already been 
studied; the others will be studied later. 

C0 2 +H 2 0^>H 2 C0 3 , carbonic acid, 

S0 3 +H 2 0->H 2 S0 4 , sulfuric acid, 
N 2 5 +H 2 0->2HN0 3 , nitric acid, 
P 2 5 H-3H 2 0->2H 3 P0 4 , phosphoric acid, 

I 2 5 +H 2 0->2HI0 3 , iodic acid. 

An oxide which by union with water forms an acid is often 
called the anhydride of the acid. 

314. Respiration. — Animals breathe air in order to obtain 
oxygen. The blood contains a complex substance, haemo- 
globin, which forms with oxygen a compound, oxyhaemoglobin, 
which easily decomposes reversibly into oxygen and haemo- 



Hydrogen and Oxygen 187 

globin. This is a typical equilibrium reaction: when oxygen, 
at the pressure at which it exists in the air, comes in contact with 
the blood in the lungs the compound is formed, 1 g. of haemo- 
globin uniting with 1.3 c.c. of oxygen; when the blood reaches 
the tissues, which take up oxygen, the compound decomposes 
and the haemoglobin is carried by the blood back to the lungs, 
where it again takes up fresh oxygen from the air. 

315. Uses of Oxygen. The Oxyacetylene Torch. — The use 
of oxygen in the oxyhydrogen blowpipe has already been men- 
tioned. By substituting acetylene for hydrogen in a blowpipe 
similarly constructed we get the oxyacetylene torch, which 
gives an intensely hot flame. It is extensively used for welding 
and for cutting iron and steel. 

Oxygen is used in several analytical processes, such as those 
studied in chapter iv. 

Deposits of carbon in the cylinders of gasoline engines are 
often removed by burning out with oxygen. Since iron burns 
also rather readily in a stream of oxygen, care must be taken to 
avoid injuring the cylinder in this way. 

316. Ozone. — When a silent electric discharge passes through 
oxygen a very remarkable change is produced ; there is a decrease 
in volume, and a gas having a powerful irritating odor is pro- 
duced. The new gas is ozone. The simplest form of apparatus 



WH 



v 




Fig. 43 

used for making ozone is shown in Fig. 43. It is a double-walled 
glass tube having the outside of the outer tube and the inside of 
the inner tube coated with tin foil. These coatings are con- 
nected by wires to the terminals of an induction coil. When the 
coil is set in action and a slow stream of oxygen is passed through 
the space between the outer and the inner tubes, the issuing gas 
is found to contain ozone. The peculiar odor of the air in the 



1 88 Introduction to General Chemistry 

neighborhood of powerful electrical machinery is due to ozone. 
Ozone is very much more active as an oxidizing agent than 
oxygen. Mercury shaken with ozone is very quickly oxidized. 
Ozone also sets iodine free from a solution of an iodide. 

If nothing but oxygen is needed to produce ozone — and such 
is actually the case — what then is the cause of the remarkable 
change in properties? In the first place it was noticed that a 
decrease in volume occurs when ozone is formed from oxygen. 
On the other hand, when ozone is changed to oxygen, as may be 
done by heating the former, the volume of the oxygen, when it 
is again cooled to the original temperature, is greater than that 
of the ozone. In fact, three volumes of oxygen give exactly two 
of ozone and vice versa. The density of ozone is one-half 
greater than that of oxygen. While 22.4 liters of oxygen weigh 
32 g., the same volume of ozone weighs 48 g. Ozone is oxygen 
in another form. If for oxygen we write the formula 2 , we must 
write 3 as the formula of ozone. The molecules of ozone 
differ from those of oxygen by containing three instead of two 
atoms. We may write the reversible equation for the relation 
between oxygen and ozone thus: 

30 2 ^20 3 . 

When ozone acts on mercury, for example, the action is as follows : 

Hg+0 3 ^HgO+0 2 . 

Only one-third of the oxygen of ozone is active, the balance changing 
into ordinary oxygen. Iodides are oxidized by ozone, thus: 

2KI+0 3 +H 2 0^2KOH+I 2 +0 2 . 

The liberated iodine may be recognized by its action on starch. 
Minute amounts of ozone may be recognized in this way, 
although the test is not conclusive proof of the presence of ozone, 
since many other substances also set iodine free from iodides. 

317. Ozone as a Germicide. — Since ozone is a very powerful 
ozidizing agent, it is not surprising that it should readily destroy 
germs. It has been found that impure water containing even 
1,000,000 bacteria per c.c. is completely sterilized by intimate 



Hydrogen and Oxygen 189 

contact with an equal volume of air containing 2 g. of ozone 
per cubic meter. In a number of important cities of Europe the 
entire municipal water supply is purified by means of ozone. 
Disinfection by chlorine is more popular. 

318. Hydrogen Peroxide, H 2 2 . — The well-known household 

antiseptic and disinfectant, hydrogen peroxide, H 2 2 , is a 3 per 

cent solution of this substance in water. Some hydrogen 

peroxide is formed by the action of sodium peroxide on ice 

water, thus: 

Na 2 2 + 2 H 2 ±5 H 2 2 + 2 NaOH. 

If water is dropped on sodium peroxide the material becomes 
very hot, and oxygen and water instead of hydrogen peroxide 
are formed (307). This is because the latter substance easily 
decomposes if hot, especially in the presence of caustic soda: 

2H 2 2 ^2H 2 0-f-0 2 . 

319. Preparation of Hydrogen Peroxide. — Barium peroxide, 
Ba0 2 , the formation of which was discussed in connection with 
Brin's process of making oxygen (310), reacts with dilute sulfuric 
acid to form hydrogen peroxide and barium sulfate : 

Ba0 2 +H 2 S0 4 ^H 2 2 +BaS0 4 . 

Since barium sulfate is insoluble in water a very pure solu- 
tion of hydrogen peroxide is easily obtained. The reaction is 
best carried out by adding finely powdered barium peroxide, 
suspended in water, very gradually to ice-cold, diluted sul- 
furic acid. The precipitate of barium sulfate is allowed to 
settle, leaving a clear solution of hydrogen peroxide. This 
solution must be made as nearly neutral as possible, other- 
wise it will decompose more or less rapidly into oxygen and 
water. 

By cautious evaporation, at a moderate temperature in a 
partial vacuum, a dilute solution of hydrogen peroxide may be 
freed from most of its water; the resulting concentrated solution 
when cooled to — io° deposits crystals of H,0 2 . 

320. Properties of Hydrogen Peroxide. — At ordinary tem- 
peratures pure hydrogen peroxide is a colorless liquid which will 



190 Introduction to General Chemistry 

mix with water in all proportions. It freezes at — 2 . It does 
not boil without decomposition, and when strongly heated it is 
liable to explode, water and oxygen being the products. The 
speed of decomposition of hydrogen peroxide at ordinary 
temperatures is greatly influenced by the presence of other sub- 
stances w T hich act as catalytic agents. Finely divided metals 
like platinum and gold cause hydrogen peroxide to decompose 
rapidly. Manganese dioxide behaves similarly. In these reac- 
tions neither the metals nor the manganese dioxide are changed. 
They are catalytic agents (303). 

Hydrogen peroxide seems to have the property of an acid, 
since it combines with some bases to form compounds which 
may be considered salts. For example, with barium hydroxide, 
Ba(0H) 2 , it reacts thus: 

H 2 2 +Ba(OH) 2 +6H 2 O^Ba0 2 -8H 2 0. 

The product consists of white crystals, rather difficultly soluble 
in water. 

Hydrogen peroxide is often used as a bleaching • agent for 
plant and animal substances, such as hair, feathers, silk, ivory, 
and straw. 

The most characteristic property of hydrogen peroxide is its 
great tendency to give up oxygen and thus to act on substances 
capable of reacting with oxygen. For example, with hydriodic 
acid it gives iodine and water: 

H 2 2 +2HI^I 2 +2H 2 0. 

The free iodine can easily be recognized by the blue color which 
it gives with a starch solution. Instead of using hydriodic 
acid we may use a solution of potassium or sodium iodide to 
which hydrochloric acid has been added, since the solution will 
then contain some hydriodic acid formed as follows : 

KI+HCteKCl+HL 

321. Detection of Hydrogen Peroxide. — A very delicate 
reaction which serves to detect small quantities of hydrogen 



Hydrogen and Oxygen 191 

peroxide is that which occurs when a solution of this substance 
is mixed with a little sulfuric acid and a very dilute solution of 
potassium dichromate, K 2 Cr 2 7 . The latter substance contains 
the element chromium, Cr, as one of its constituents. The 
solution turns blue, and when a little ether is added and shaken 
up with the blue solution the ether dissolves the blue substance. 
If the mixture is allowed to stand a minute or two the blue ether 
solution separates from the water solution, on which it floats as a 
blue layer. 

Other reactions of hydrogen peroxide are discussed in the 
following chapter (347, 348). 

322. Peroxides and Dioxides. — We have just learned that 
hydrogen peroxide is formed by the action of dilute acids on 
Na 2 2 and Ba0 2 , and we might therefore be inclined to expect 
that we should also get H 2 2 by the action of acids on Pb0 2 and 
Mn0 2 . But this is not the case; no H 2 2 can be obtained in 
any way from these last-mentioned oxides. For this reason 
these oxides of lead and manganese are called dioxides to dis- 
tinguish them as a class from those which yield H 2 2 and which 
are called peroxides. Thus we call Ba0 2 barium peroxide and 
Pb0 2 lead dioxide. 

323. Graphic Formulae. — We have so far considered that 
oxygen has a valence (183) of two, or is bivalent, since in water 
two symbol weights of hydrogen are united with one of oxygen. 
But what then is the valence of oxygen in H 2 2 ? In order to be 
able to answer this question, we must consider the matter of 
valence from the standpoint of the atomic-molecular hypothesis. 
We have learned (221) that the molecule of water is made up 
of two atoms of hydrogen and one of oxygen. Since all the 
molecules of water are made up in just this fashion, it would seem 
to follow that the three atoms must be related to one another 
in some very definite way. We may think of them as being 
joined to one another, in which case there are the two possibilities 
indicated by the following graphic formulae in which the lines 
joining the symbols are called bonds. 

(1) H-H-0 (2) H-O-H. 



192 Introduction to General Chemistry 

The first graphic formula indicates that one of the hydrogen 
atoms is attached on the one hand to the atom of oxygen and 
on the other to the second atom of hydrogen. Formula (2) indi- 
cates that it is the oxygen atom which is attached on either hand 
to an atom hydrogen. It is obvious that the second formula 
is the more consistent, since in it both atoms of hydrogen are 
attached by single bonds to the atom of oxygen which holds an 
atom of hydrogen by each of its two bonds. In formula (1) the 
middle hydrogen atom is represented as having two bonds, while 
the other hydrogen atom and also the atom of oxygen are shown 
as having but one bond each. Since we think of all atoms of 
hydrogen as being alike, we must reject the first formula in favor 
of the second. 

When viewed in the above-mentioned manner the valence of 
an element is seen to be the holding capacity of its atoms for atoms 
of hydrogen or other univalent elements like chlorine. 'We may 
therefore think of an atom of oxygen which is bivalent as having 
two valence bonds, each of which can hold one atom of a univalent 
element. 

324. The Graphic Formulae of Peroxides. — We are now 
prepared to consider the question of the valence of oxygen in 
H 2 2 and the graphic formula of this substance. At the outset 
it may be stated that the valence of hydrogen is considered by 
chemists to be invariably one. If the valence of oxygen is taken 
to be two, there is but one possible graphic formula, namely: 

H-O-O-H. 

This is the commonly accepted formula. For sodium peroxide 
we then have the formula, 

Na-O-O-Na, 
and for barium peroxide, 

Ba< I 



Hydrogen and Oxygen 193 

In their dioxides, manganese and lead have without doubt 
a valence of four, or are tetravalent, and oxygen is, as usual, 
bivalent. We therefore write for these oxides the formulae 

= Mn=0 and = Pb = 0, 

thereby indicating that each oxygen atom is attached to an 
atom of manganese or lead by two bonds, or in other words by a 
double bond. 

The monoxides of manganese, MnO, and lead, PbO, have 
their atoms doubly bound, thus : 

Mn=OandPb = 0. 
In these oxides both metals are bivalent. 



CHAPTER XV 
OXIDATION AND REDUCTION 

325. Oxidation. — When a substance unites with oxygen it 
is said to be oxidized. Hydrogen when burned is oxidized, giv- 
ing water. Metals are said to be oxidized when they combine 
with oxygen; for example, 

* 2Cu+0 2 ->2CuO. 

By certain indirect methods a lower oxide of copper, cuprous 
oxide, Cu 2 0, can be made. This oxide can unite with more 
oxygen if heated in air or in oxygen and form the common oxide, 
CuO, which is known also as cupric oxide, in order to distinguish 
it from the lower oxide: 

2Cu 2 0+0 2 ->4CuO. 

We say in this case that cuprous oxide has been oxidized to cupric 
oxide. 

The action of oxygen gas on hydrogen chloride at a high 
temperature (239) proceeds according the equation 

4 HCl+0 2 ->2Cl 2 +2H 2 0, 

and in consequence we say that the hydrogen chloride has been 
oxidized. 

326. Oxidizing Agents. — Very often substances may be 
oxidized by compounds of oxygen as well as by oxygen itself. 
For example, heated copper oxide, CuO, oxidizes hydrogen and 
all its compounds, such as ammonia, acetylene, etc. We say, 
therefore, that copper oxide is an oxidizing agent. Any sub- 
stance which oxidizes another is called an oxidizing agent. Lead 
dioxide and manganese dioxide are powerful oxidizing agents, 
as shown by the fact that each is able to oxidize hydrochloric 

acid, 

MnO a +4HCl->MnCl a +Cl a +2H 2 0. (234) 

194 



Oxidation and Reduction 195 

Potassium permanganate, KMn0 4 , which also easily oxi- 
dizes hydrochloric acid, is one of the most powerful of all oxidiz- 
ing agents. It acts according to the equation 

2 KMn0 4 + i6HCl-> 2 KCl+ 2 MnCl 2 + 5 C1 2 -J-8H 2 0. (235) 

Nitric acid and its salts (104), the nitrates, are good oxidiz- 
ing agents. Gunpowder is a mixture of finely powdered potas- 
sium nitrate, charcoal, and sulfur. The explosion of gunpowder 
is due to the extremely rapid oxidation of the charcoal (carbon) 
and sulfur to carbon dioxide, C0 2 , and sulfur dioxide, S0 2 , the 
oxygen being furnished by the potassium nitrate, KN0 3 . The 
other products of the explosion are nitrogen and potassium sul- 
fide, K 2 S. 

Potassium chlorate, KC10 3 , is a powerful oxidizing agent, 
readily giving up all its oxygen to oxidizable substances and 
leaving the chloride, KC1. 

Sulfuric acid (93) is capable of oxidizing some substances; 
hot sulfuric acid acts on charcoal thus: 

2 H 2 S0 4 + C->C0 2 + 2S0 2 + 2H 2 0. 

It is probable that each molecule of sulfuric acid first loses one 
atom of oxygen giving a molecule of sulfurous acid, H 2 S0 3 , and 
that this substance, which is unstable, then decomposes into 
sulfur dioxide and water: 

H 2 S0 3 ->S0 2 +H 2 0. (286) 

327. Reduction and Reducing Agents. — When hydrogen is 
oxidized by hot copper oxide, thus: 

CuO+H 2 ->Cu+H 2 0, 

the copper oxide is said to be reduced to metallic copper. In 
consequence we call hydrogen a reducing agent. Any substance, 
as, for example, acetylene or methane, which can reduce copper 
oxide is also called a reducing agent. In any reaction, if one 
substance is oxidized the oxidizing agent is by necessity reduced; 
oxidation and reduction always go on together. All substances 
which are acted upon by oxidizing agents are, of course, redu- 
cing agents. 



196 Introduction to General Chemistry 

328. Carbon as a Reducing Agent. — Since charcoal, which 
is nearly pure carbon, burns readily, it is capable of taking up 
oxygen from oxidizing agents and is therefore a good reducing 
agent. A mixture of powdered copper oxide and charcoal reacts 
vigorously, if strongly heated, giving copper and carbon dioxide: 

2CuO+C->2Cu+C0 2 . 

In this reaction the copper oxide is the oxidizing agent and the 
charcoal (carbon) the reducing agent. 

Many other metallic oxides can be reduced in a similar man- 
ner by carbon. In place of charcoal, coke or coal, which are 
largely carbon, may be used. Thus ferric oxide, Fe 2 3 , which 
in the form of the mineral hematite is the most important ore 
of iron, is reduced by coke at a white heat to metallic iron. The 
result may be represented by the equation 

2Fe 2 3 +3C->4Fe+ 3 C0 2> 

although it is very probable that the reaction is less simple under 
the conditions actually met with in practice. 

329. Carbon Monoxide as a Reducing Agent. — When carbon 
is burned in a deficient supply of air, carbon monoxide, CO, is 
formed instead of dioxide : 

2C + 2 ->2CO. 

This is a colorless, odorless, and very poisonous gas which will 
burn with a nearly non-luminous flame to form carbon dioxide, 

2CO+0 2 ^2C0 2 . 

Some oxidizing agents are able to oxidize carbon only to 
monoxide and not to dioxide. Zinc oxide behaves in this way: 

ZnO+C^Zn+CO. 

This is the reaction by which zinc is made from its ores. The 

reaction between ferric oxide and carbon can also give carbon 

monoxide, 

Fe 2 3 + 3 C->2Fe+ 3 CO. 

But carbon monoxide can also reduce ferric oxide: 
Fe 2 3 -f 3 CO->2Fe+ 3 C0 2 . 



Oxidation and Reduction 197 

The last two equations doubtless represent the steps by which 
ferric oxide and carbon react to give iron and carbon dioxide 
(328). 

330. Aluminum as a Reducing Agent. — Metallic aluminum 
unites vigorously with oxygen at a white heat, although it has 
no tendency to oxidize in the air at ordinary temperatures. The 
burning of aluminum occurs thus: 

4Al+30 2 ->2Al 2 3 - 

When a mixture of powdered aluminum and ferric oxide is 
strongly heated a very violent reaction takes place, giving iron 
and aluminum oxide: 

2Al+Fe 2 3 ->2Fe+Al 2 3 . 

The mixture of aluminum and ferric oxide has been given the 
trade name of thermite by its inventor, Goldschmidt, who uses 
it to make small quantities of molten iron for the repair of broken 
iron castings, etc. 

Many other metallic oxides can also be reduced by aluminum. 

331. Oxidation Considered as a Change of Valence. — We 
have already learned (173) that iron forms two series of com- 
pounds , ferrous and ferric, as illustrated by the following formulae : 



Ferrous Compounds 


Ferric Compound 


FeO 


Fe 2 3 


Fe(OH) 2 


Fe(OH) 3 


FeCL 


FeCl 3 


FeBr 2 . 


FeBr 3 


Fe(N0 3 ) 2 


Fe(N0 3 ) 3 


FeS0 4 


Fe 2 (S0 4 ) 3 



The valence of iron is two in ferrous compounds and three in 
ferric. According to this usage of the term valence we should 
be forced to say that the valence of free or uncombined iron is 
zero. 

If free iron is changed into ferrous oxide, 

2Fe+0 2 ->2FeO, 



198 Introduction to General Chemistry 

it is oxidized, and its valence is increased from zero to two. 
Moreover, if ferrous oxide is changed into ferric oxide, 

4FeO+0 2 ^2Fe 2 3 , 

it is also plain that the iron is further oxidized, and that its 
valence has increased from two to three. It is customary to 
say that ferric oxide is a higher oxide of iron than ferrous oxide ; 
or that iron in ferric oxide is in a higher state of oxidation than 
in ferrous oxide. 

In the case of iron and its oxides we see that the oxidation of 
iron and the increase in its valence go hand in hand. 

With respect to other elements that unite, with oxygen we 
also find that their oxidation results in an increase in their 
valence. A few additional examples will help to illustrate this 
point. In the change of copper into cuprous oxide, Cu 2 (325), 
the oxidation of the copper is accompanied by an increase of its 
valence from zero to one. In cuprous oxide copper is univalent 
(146). In the oxidation of cuprous oxide to cupric oxide, 

2Cu 2 0+0 2 -> 4 CuO, (325) 

the valence of copper is increased from one to two. In cupric 
oxide copper is bivalent (146). 

When carbon is oxidized to carbon monoxide, 

2C+0 2 ->2CO, (329) 

the valence of carbon is increased from zero to two (carbon is 
bivalent in carbon monoxide). In the oxidation of carbon 
monoxide to carbon dioxide, 

2 CO+0 2 ->2C0 2 , (329) 

the valence of carbon is increased to four, carbon becoming 
quadrivalent. 

332. A Broader Meaning of the Term Oxidation. — Since in 
the change of any ferrous compound into the corresponding 
ferric compound (173, 331) the valence of iron always increases 
from two to three, all such changes may well be considered to 
be of the same class. It has become the custom among chemists 



Oxidation and Reduction 199 

to call such increase of valence of iron an oxidation of the iron 
irrespective of the nature of the element or radical combined 
with the iron. Thus in the reaction 

2FeCl 2 +Cl 2 ->2FeCl 3 , 

whereby ferrous chloride is changed to ferric chloride, we say 
that the iron has been oxidized. The only question that should 
arise here is : Why call this increase in valence of iron an oxida- 
tion in cases where no oxygen is involved? We can only say 
that it is a custom sanctioned by long and universal usage. 

By way of further illustration of the use of the term oxida- 
tion in its broader sense we may cite the following examples. 
When metallic sodium is changed into chloride, NaCl, or nitrate, 
NaN0 3 , its valence is increased from zero to one, and we say 
that the sodium has been oxidized. When zinc is changed into 
oxide, ZnO; sulfate, ZnS0 4 ; chloride, ZnCl 2 ; or nitrate, Zn(N0 3 ) 2 
(148), we say that the zinc has undergone oxidation; and further- 
more, since in all these compounds zinc is bivalent, we say that 
zinc in all these compounds is in the same state or stage of oxida- 
tion. In fact, zinc in its compounds is always bivalent. 

333. Review of Other Elements with Variable Valence. — 
Iron is not the only element having a variable valence. We 
have already seen (179, 180) that mercury also forms two series 
of compounds, the mercurous, in which the element has a valence 
of one, and the mercuric, where the valence is two, as illustrated 
by the following formulae: 



Mercurous Compounds 


Mercuric Compounds 


Hg 2 


HgO 


HgCI 


HgCI 3 


Hgl 


Hgl 2 


HgN0 3 


Hg(N0 3 ) 2 


Hg 2 S0 4 


HgS0 4 



Mercurous compounds are converted into mercuric by oxida- 
tion, and mercuric into mercurous by reduction. 

Copper also forms two series of compounds, cuprous and 
cupric. We know cuprous oxide, Cu 2 0, and cuprous chloride, 
CuCl, as well as the commoner cupric compounds, such as cupric 



200 Introduction to General Chemistry 

oxide, CuO, cupric chloride, CuCl 2 , cupric sulfate, CuS0 4 , 
etc. (165). 

It will be noted that compounds representing the lower state 
of oxidation have names ending in ous, while those corresponding 
to the higher state of oxidation end in ic. 

334. Another Class of Oxidation Reactions. — We have 
given as one illustration of an oxidation reaction the action of 
manganese dioxide on hydrochloric acid: 

Mn0 2 + 4 HCl->MnCl 2 +Cl 2 +2H 2 (326) 

In this reaction the manganese dioxide is the oxidizing agent. 
The chlorine of the hydrochloric acid in being set free has its 
valence decreased from one to zero. The valence of the hydro- 
gen remains unchanged in the reaction: it is neither oxidized 
nor reduced. Now we see that when combined chlorine, as in a 
chloride, is set free its valence is decreased, and that in this change 
the chlorine is oxidized. In the case of metallic elements oxida- 
tion involves increase of valence (331). We see, therefore, that 
in the case of chlorine, a non-metallic element, its oxidation in 
cases like that cited involves a decrease in its valence. Other 
non-metallic elements, like bromine, iodine, and sulfur, when set 
free by oxidizing agents from their compounds with hydrogen 
or metals are oxidized, while at the same time there occurs a 
decrease in valence. 

335. Reduction and Change of Valence. — That valence 
changes accompany reduction, as well as oxidation, will be at 
once apparent by the consideration of any reaction in which 
reduction takes place. Take, for example, the simple case of 
the reduction of cupric oxide by hydrogen, 

CuO+H 2 ->Cu+H 2 0. (327) 

The copper is reduced to the free state, its valence changing 
from two to zero, while the oxygen merely changes partners 
without change of valence and is therefore neither oxidized nor 
reduced. When any oxide of a metal is reduced by hydrogen, 
carbon, or any other reducing agent, the valence of the metal is also 
reduced or lowered. 



Oxidation and Reduction 



201 



On the other hand, in the reaction 
C1 2 +H 2 -> 2 HC1 

we see that the reduction of the non-metallic element chlorine 
is accompanied by an increase in its valence from zero to one. 
In general, when a non-metallic element like a halogen, sulfur, or 
oxygen itself unites with hydrogen or a metal, the non-metal is 
reduced, while concurrently its valence is increased. 

336. Classification of Valence Changes. — In analyzing the 
foregoing cases we have seen that there is a difference in 
the behavior of metals and non-metals. If we consider the 
many compounds which we have studied, we shall see that 
metals and hydrogen unite with non-metals more or less 
readily but do not unite with each other, at least in the cases 
studied. On the other hand non-metals not only unite with 
metals but also unite among themselves, as in the case of 
carbon and oxygen or sulfur and oxygen (286). 

In all but the rarest cases the valence of an atom in com- 
bination may be found directly or indirectly from the following 
rules, which are established by experience: hydrogen (sodium 
and potassium) are univalent in all their compounds; oxygen 
is bivalent in all its compounds except the peroxides. The 
latter have definite characteristics and may be identified. The 
classification of the changes of valence may be made systemati- 
cally on the basis of the change of the valence of an atom toward 
a metal or non-metal as indicated in the table below. Whether 
the atom in question is a metal or a non-metal does not matter. 



For the Atom in 
Question 



Valence toward 
non-metals . . . 

Valence toward 
metals or hy- 
drogen 



In Oxidation 



In Reduction 



Increases 



Decreases 



Decreases 



Increases 



If the reactions of reagents are classified according to this scheme 
those which are found to undergo oxidation or reduction, and 



202 Introduction to General Chemistry 

so are themselves either reducing agents or oxidizing agents, 
will be found actually to have similar chemical activities. 

337. Two Important Kinds of Reactions. — But very few of 

the reactions of substances in solution studied in chapters prior 

to the present involve oxidation and reduction. In such 

reactions as 

HCl+NaOH->NaCl+H 2 0, 

AgN0 3 +HCl->AgCl+HN0 3 , (169) 

and 

FeCl 3 +3NaOH->Fe(OH) 3 +3NaCl, (173) 

neither oxidation nor reduction occurs. These are called double- 
decomposition reactions. In such reactions no element changes 
its valence. 

If oxidation and reduction take place, the reaction is of a 
distinctly different kind. In such reactions two or more elements 
change valence. 

338. Intensity of Activity of Oxidizing and Reducing Agents. 
— Both oxidizing and reducing agents differ greatly in their 
intensity of activity. For example, manganese dioxide will 
oxidize cold, dilute hydrochloric acid, but oxygen gas will not. 
In consequence we say that manganese dioxide is a stronger 
or more powerful oxidizing agent than oxygen itself. When we 
rind, as we may readily do by experiment, that dilute nitric acid 
will oxidize a ferrous salt to a ferric salt and that dilute sulfuric 
acid will not do so, we conclude that nitric acid is a stronger or 
better oxidizing agent than sulfuric acid. When we know that 
hydriodic acid will reduce sulfuric acid, but that hydrochloric 
acid will not do so, we conclude that hydriodic acid is a better 
reducing agent than hydrochloric acid. 

339. Hydrogen Sulfide, H 2 S. — Iron and sulfur unite directly 
at a red heat to form ferrous sulfide, FeS : 

Fe+S^FeS. 

This is a black solid, which is insoluble in water. It reacts 
readily with hydrochloric acid to give ferrous chloride and 
hydrogen sulfide: 

FeS+2HCl->FeCl 2 +H 2 S. 



Oxidation and Reduction 203 

Hydrogen sulfide is a colorless gas, of which water dissolves 

three to four times its own volume. It has a very disagreeable 

odor, resembling rotten eggs, and is extremely poisonous. 

Fatal accidents have often occurred from breathing the gas. 

Hydrogen sulfide is a very powerful reducing agent. In water 

solution it is easily oxidized by atmospheric oxygen, giving sulfur 

and water : 

2H 2 S+0 2 ->2S+2H 2 0. 

A water solution of hydrogen sulfide reacts rapidly with iodine 
to form hydriodic acid (265) and sulfur: 

H 2 S+I 2 ->2HI+S. 

This reaction furnishes a very good practical method for making 
hydriodic acid. We have only to pass hydrogen sulfide gas into 
water containing powdered iodine. When all the iodine has 
been reduced, the solid sulfur can be filtered out, giving a 
clear, colorless filtrate which contains only hydriodic acid and 
water. 

Hydrogen sulfide readily reduces dilute sulfuric acid, which 
is but a very mild oxidizing agent capable of oxidizing only the 
most active reducing agents; the products are sulfurous acid, 
H 2 S0 3 , sulfur, and water: 

H 2 S+H 2 S0 4 ->H 2 S0 3 +S+H 2 0. 

Hydrogen sulfide can be made by the action of sulfuric acid 
on ferrous sulfide, thus: 

FeS+H 2 S0 4 ->FeS0 4 +H 2 S; 

but this is not advisable in practice because of the interaction of 
the sulfuric acid with the hydrogen sulfide. Hydrochloric acid 
is the best acid to use in making hydrogen sulfide. 

Hydrogen sulfide is oxidized by all except the very mildest 
oxidizing agents. As a final example of its behavior, its action 
on ferric salts may be given. These are reduced to ferrous salts, 
thus: 

2FeCl 3 +H 2 S^2FeCl 2 +S+2HCl. 



204 Introduction to General Chemistry 

340. Sulfurous Acid, H 2 S0 3 . — When sulfur burns, it forms 
sulfur dioxide, S0 2 , a colorless gas, with a strong odor: 

S+0 2 ->S0 2 . (286) 

Sulfur dioxide is very soluble in water, with which it unites 
partially to form sulfurous acid: 

S0 2 +H 2 O^H 2 S0 3 . (286) 

This reaction is reversible; by boiling a solution of sulfurous 
acid the latter can be completely decomposed and all sulfur 
dioxide driven off. 

Sulfurous acid is a reducing agent which, when oxidized, is 
converted into sulfuric acid. It reacts slowly with atmospheric 
oxygen, thus: 

2H 2 S0 3 +0 2 ->2H 2 S0 4 . 

It is rapidly oxidized by chlorine, which is thereby reduced 
to hydrochloric acid: 

H 2 S0 3 +C1 2 +H 2 0->H 2 S0 4 +2HC1. 

Manganese dioxide and sulfurous acid react as follows : 
Mn0 2 +H 2 S0 3 ->MnS0 4 +H 2 0. 

Ferric salts are reduced to ferrous salts by sulfurous acid, as 
illustrated by the following equation: 

Fe 2 (S0 4 ) 3 +H 2 S0 3 +H 2 0->2FeS0 4 +2H 2 S0 4 . 

341. Hydrochloric, Hydrobromic, and Hydriodic Acids as 
Reducing Agents. — These acids in water solution can all be 
oxidized, and are therefore to be considered as reducing agents. 
Hydriodic acid is the most easily oxidized of the three, and is 
therefore the best or most powerful reducing agent. It is 
oxidized by atmospheric oxygen (265), which has no action 
whatever on a water solution of hydrochloric acid. The latter 
substance acts as a reducing agent only with respect to the most 
powerful oxidizing agents, such as manganese (234) and lead 
dioxides (235) and potassium permanganate (236). Hydro- 



Oxidation and Reduction 205 

bromic acid is a better reducing agent than hydrochloric acid / 
but not as powerful as hydriodic acid. 

Hydrobromic acid, as a reducing agent, reacts with concen- 
trated sulfuric acid, as an oxidizing agent, as follows, 

2HBr+H 2 S0 4 ->Br 2 +S0 2 +2H 2 0, 

forming free bromine, sulfur dioxide, and water. 

Hydriodic acid reacts even more vigorously with concentrated 
sulfuric acid. In this case the products vary according to the 
proportions taken, but the reduction of the sulfuric acid may go 
as far as the formation of free sulfur and hydrogen sulfide. The 
possible reactions are represented in the following equations : 

H 2 S0 4 + 2ffl->I 2 -f-S0 2 + 2 H 2 0, 

H 2 S0 4 +6HI-> 3 I 2 +S+ 4 H 2 0, 

H 2 S0 4 +8HI-> 4 I 2 +H 2 S+ 4 H 2 0. 

It will now be understood why roundabout methods are used 
to prepare hydrogen bromide and hydrogen iodide instead of 
the simple reaction with concentrated sulfuric acid and a salt, 
as is done in the preparation of hydrogen chloride. 

342. Manganese and Its Compounds. — Manganese, Mn, is a 
metallic element which in the free form resembles iron rather 
closely. Its principal ore is the dioxide, Mn0 2 , called by 
mineralogists pyrolusite. Manganese forms a series of salts 
corresponding to the salts of ferrous iron; among such we have 
manganous chloride, MnCl 2 , manganous nitrate, Mn(N0 3 ) 2 , and 
manganous sulfate, MnS0 4 . These salts are pale pink in color 
and are easily soluble in water. The dilute solutions, which are 
almost colorless, give with sodium hydroxide white precipitates 
of manganous hydroxide: 

MnCl 2 + 2NaOH->Mn(OH) 2 + 2 NaCl. 

This hydroxide corresponds to an oxide, MnO: 
Mn(OH) 2 ->MnO+H 2 0. 

In all these compounds except the dioxide, Mn0 2 , manganese is 
bivalent; in the dioxide it is quadrivalent. 



206 Introduction to General Chemistry 

Manganese forms a variety of compounds of a very different 
character from the ones just mentioned; of these the most 
important is potassium permanganate. 

343. Potassium Permanganate, KMn0 4 .— This substance is 
the potassium salt of permanganic acid, HMn0 4 , in which 
manganese acts as an acid-forming element. The salt is made 
from manganese dioxide and potassium hydroxide by compli- 
cated reactions which need not be considered at present. It 
forms dark-purple crystals which dissolve in water to form a 
purple solution having nearly the color of the vapor of iodine. 
It is a very important substance and is one of the most powerful 
of all oxidizing agents. 

We have already learned that potassium permanganate 
oxidizes hydrochloric acid, thus : 

2 KMn0 4 +i6HCl->2KCl+2MnCl 2 +5Cl 2 +8H 2 0. (235) 

It can also .oxidize almost any substance which is capable of 
being oxidized in solution. Two additional examples may be 
given as illustrations : 

2KMn0 4 +5H 2 S0 3 ->K 2 S0 4 +2Mnsb 4 +2H 2 S0 4 +3H 2 0, 
2KMn0 4 +ioFeS0 4 +8H 2 S0 4 ->K 2 S0 4 +2MnS0 4 +5Fe 2 (S0 4 ) 3 +8H 2 0. 

In the last reaction the sulfuric acid acts neither as a reducing 
nor an oxidizing agent, but is used to keep the solution acid. The 
sulfate radical is not decomposed in the reaction. 

From the foregoing equations it is apparent that when two 
molecules of permanganate change to manganese sulfate or 
chloride, they change the valence of ten atoms (chlorine in 
hydrochloric acid or iron in a ferrous salt) by one unit of valence 
each, or of five atoms by two units of valence each (sulfur in 
sulfurous acid). This relationship exists because the valence 
of manganese, which is seven in permanganate, changes to two 
in manganese sulfate or chloride. 

344. Chromium and Its Compounds. — The element chro- 
mium, Cr, is a hard metal, resembling iron in appearance. It 
forms a series of salts of which chromic chloride, CrCl 3 , and 
chromic sulfate, Cr 2 (S0 4 ) 3 , are typical examples. Solutions of 



Oxidation and Reduction 207* 

chromic salts are either green or violet in color, according to the 
method of preparation. These solutions give with ammonium 
hydroxide bluish precipitates of chromic hydroxide : 

CrCl 3 +3NH 4 OH^Cr(OH) 3 + 3 NH 4 Cl. 

The hydroxide when strongly heated gives chromic oxide: 
2Cr(OH) 3 ->Cr 2 3 +3H 2 0. 

It will be seen that chromic salts are analogous to ferric salts 
and that in these compounds chromium is trivalent. 

345. Chromates and Dichromates. — When chromic oxide is 
fused with sodium nitrate or sodium peroxide, sodium chromate, 
Na 2 Cr0 4 , is formed. This is a bright-yellow crystalline salt, 
readily soluble in water. It may be considered as derived from 
chromic acid, H 2 Cr0 4 . Potassium chromate, K 2 Cr0 4 , is also a 
yellow crystalline salt which is readily made by methods similar 
to those that give the sodium salt. 

A solution of potassium chromate, which is bright yellow in 
color, turns deep orange when mixed with sulfuric acid. The 
solution contains potassium dichromate, K 2 Cr 2 7 , which has 
been formed thus: 

2 K 2 Cr0 4 +H 2 S0 4 ->K 2 Cr 2 7 +K 2 S0 4 +H 2 0. 

Potassium dichromate forms orange-colored crystals, which 
dissolve in water to form an orange-colored solution. In the 
foregoing reaction we might have expected to get potassium 
hydrogen chromate, KHCr0 4 ; but if this salt is first formed it 
decomposes at once, as follows : 

2 KHCr0 4 ->K 2 Cr 2 7 +H 2 0. 

Sodium dichromate, Na 2 Cr 2 7 , orange-colored crystals, can 
be made in a similar manner from sodium chromate. 

346. Chromates and Dichromates as Oxidizing Agents. — 
Solutions of either chromates or dichromates are oxidizing 
agents. More commonly a strongly acid solution is used. Thus 
with hydrogen sulfide an acid solution of potassium chromate 



•2o8 Introduction to General Chemistry 

(potassium dichromate) reacts to form sulfur and chromium 
salts : 

2K 2 Cr0 4 +3H 2 S+ioHC1^4KCl+2CrCl 3 +8H 2 O+ 3 S; 

or if the equation is written for the dichromate we have 
K 2 Cr 2 7 +3H 2 S+8HCl->2KCl+2CrCl 3 +7H 2 0+3S. 

With sulfurous acid as a reducing agent the reaction yields 
sulfuric acid and chromium sulfate : 

2K 2 Cr0 4 +3H 2 S0 3 +2H 2 S0 4 ->2K 2 S0 4 +Cr 2 (S0 4 )3+5H 2 0. 

Apparently two molecules of potassium chromate (or one of the 
dichromate) can cause six units of valence change on other 
atoms, three of sulfur in H 2 S if free sulfur is the product, or 
three of sulfurous acid to form sulfuric acid. 

It is plain that permanganates are more powerful oxidizers 
than chromates or dichromates, since the first can oxidize 
hydrochloric acid and the second cannot, except in very con- 
centrated acid solution. 

347. Hydrogen Peroxide as an Oxidizing Agent. — We have 

already learned (318) that hydrogen peroxide easily decomposes 

into water and oxygen, and that for this reason it acts as an 

oxidizing agent. Its action on hydriodic acid was shown to take 

place thus: 

H 2 2 + 2HI^I 2 + 2 H 2 0. (320) 

Sulfurous acid is readily oxidized to sulfuric acid : 

H 2 2 +H 2 S0 3 ->H 2 S0 4 +H 2 0. 

Ferrous salts are oxidized to ferric salts as the following 
equation will illustrate : 

H 2 2 +2FeS0 4 +H 2 S0 4 ->Fe 2 (S0 4 ) 3 +2H 2 0. 

Lead forms with sulfur lead sulfide, PbS, a black substance, 
almost insoluble in water. It is obtained as a black precipitate 
by the action of hydrogen sulfide on a solution of a lead salt: 

Pb(N0 3 ) 2 +H 2 S->PbS+2HN0 3 . 



Oxidation and Reduction 209 

Hydrogen peroxide oxidizes lead sulfide to lead sulfate (167) : 
PbS+4H 2 2 ->PbS0 4 + 4 H 2 0. 

Since lead sulfate is white, the effect of the action is easily seen. 
The blackening of old oil paintings is due to the gradual conver- 
sion of the lead compounds that have served as ingredients of 
the paint into lead sulfide by the action of sulfur compounds 
occurring in the air. Blackened paintings are often restored 
to their original colors by treating them with hydrogen peroxide, 
which converts the black lead sulfide into white lead sulfate. 

It has already been mentioned that animal and vegetable 
substances are bleached by hydrogen peroxide. The exact 
nature of the changes that occur in such reactions is not in general 
known, but it is safe to conclude that they are processes of oxida- 
tion which convert colored into colorless substances. 

348. The Reducing Action of Hydrogen Peroxide. — The 
action of hydrogen peroxide on silver oxide yields free silver, 
and we may say that the silver oxide has been reduced. 

H 2 2 +Ag 2 0-> 2Ag+H 2 0+0 2 . 

Another important reaction of this class is found in the action 
of hydrogen peroxide on potassium permanganate in acid solu- 
tion, which takes place thus: 

5 H 2 2 + 2KMn0 4 +3H 2 S0 4 ^> K 2 S0 4 + 2 MnS0 4 + 8H 2 0+ 50 2 . 

The products are the colorless solution of the sulfates of potas- 
sium and manganese in addition to free oxygen. 

349. Hypochlorous Acid, HCIO. — It is probable that chlorine 
reacts reversibly with water in which it is dissolved to form 
hydrochloric acid and hypochlorous acid, HCIO, thus: 

C1 2 +H 2 0^HC1+HC10. 

Since this is a reversible reaction, all four substances are con- 
tained in equilibrium in a solution of chlorine in water. Hypo- 
chlorous acid is very unstable, that is, it easily decomposes. 
and for this reason it cannot be obtained except in the form of a 
dilute water solution. It has only very weak acid properties 



210 Introduction to General Chemistry 

and cannot even decompose calcium carbonate, which is acted 
upon by almost all other acids. As is well known, hydrochloric 
acid reacts with calcium carbonate as follows: 

2HCl+CaC0 3 ->CaCl 2 +C0 2 +H 2 0. (163) 

As a matter of fact, when calcium carbonate is added to chlorine 
water it reacts as follows : 

CaC0 3 +2Cl 2 +H 2 0-> 2HC10+CaCl 2 +C0 2 . 

From the resulting solution hypochlorous acid mixed with much 
water vapor can be driven off by cautious heating; the condensed 
vapor forms a dilute solution of hypochlorous acid. This reac- 
tion seems to prove that chlorine and water react to form hydro- 
chloric and hypochlorous acids. 

350. Hypochlorites. — If chlorine is passed into a cold, dilute 
solution of sodium hydroxide, sodium chloride and sodium hypo- 
chlorite, NaCIO, are formed: 

Cl 2 +2NaOH->NaCl+NaC10+H 2 0. 

This is exactly what we should expect if both acids which result 

from the action of chlorine on water are neutralized by the 

sodium hydroxide. Chlorine and potassium hydroxide react 

similarly : 

Cl 2 +2KOH->KCl+KC10+H 2 0. 

351. Bleaching Powder. — The action of chlorine gas on solid 
slaked lime, calcium hydroxide, takes place thus: 

2Cl 2 +2Ca(OH) 2 ->CaCl 2 +Ca(C10) 2 +2H 2 0. 

The product of the reaction is a white powder known as chloride 
of lime or bleaching powder. It is a mixture of calcium chloride 
and calcium hypochlorite. It is extensively used in the bleaching 
of cotton goods and for a variety of other purposes. Before 
taking up the chemical behavior of hypochlorous acid and hypo- 
chlorites it will be of interest to consider the formation of these 
substances from the standpoint of oxidation and reduction. 



Oxidation and Reduction 211 

352. The Oxidation Products of Chlorine. — By the action of 
chlorine gas on dry mercuric oxide, HgO, chlorine monoxide, 

C1 2 0, a brownish-yellow gas, is obtained. It is obvious that in 
this reaction the chlorine has been oxidized. Now this oxide 
of chlorine unites with water to form hypochlorous acid, 

C1 2 0+H 2 0->2HC10. 

The relation between chlorine monoxide and hypochlorous 
acid is similar to that between sulfur dioxide and sulf urous acid : 

S0 2 +H 2 0->H 2 S0 3 . 

Chlorine and sulfur also show similar behavior in that each forms 
compounds with hydrogen and metals, namely chlorides and 
sulfides. 

353. The Formation of Chlorates. — Hypochlorites are very 
unstable salts. A warm, concentrated solution of sodium hypo- 
chlorate changes more or less rapidly into sodium chloride and 
sodium chlorate, NaC10 3 , according to the equation, 

3 NaC10-> 2NaCl+NaC10 3 . 

Potassium hypochlorite changes in a similar fashion, yielding 
potassium chlorate, KC10 3 . 

Sodium and potassium chlorates are powerful oxidizing 
agents, since they contain large proportions of easily liberated 
oxygen. When the dry crystals are heated they decompose 
finally into chlorides and oxygen: 

2 KC10 3 ->2KCl+30 2 . 

This reaction takes place in two stages (306). The first change 
gives rise to a perchlorate, KC10 4 , thus: 

ioKC10 3 ->4KCl+6KC10 4 + 3 2 . 

Until recently potassium chlorate was used extensively, and 
sodium chlorate was rarely seen. The reason was twofold: in 
the first place potassium chlorate was made very largely in 
Germany, where potassium compounds are cheap on account 
of the immense potash deposits found in that country; and in 



212 Introduction to General Chemistry 

the second place sodium chlorate, being more soluble, is more 
difficultly purified than the potassium salt. Since the war began 
there has been a shortage of potash, because no other country 
besides Germany has much easily accessible potash. As a 
consequence the manufacture of sodium salts has been stimu- 
lated, and since 191 5 there has been an abundant supply of 
sodium chlorate. This can be used advantageously in place of 
potassium chlorate for nearly all purposes. 

354. Chloric Acid and Chlorine Dioxide. — Potassium and 
sodium chlorates are salts of chloric acid, HC10 3 . This is a very 
unstable acid, which is known only in dilute solution. Upon 
evaporation of the solution the acid decomposes, giving chlorine 
dioxide, C10 2 , and other products. 

If a few drops of concentrated sulphuric acid are poured on 
a small crystal of sodium chlorate in a dry test tube, a yellow 
gas forms, which explodes with violence a few seconds later. 
This dangerous experiment should be performed with great 
caution. The yellow gas is chlorine dioxide, C10 2 , which was 
formed by the decomposition of the chloric acid set free, thus : 

NaC10 3 +H 2 S0 4 ->NaHS0 4 +HC10 3 . 

The explosion of chlorine dioxide is due to decomposition into 
its elements: 

2C10 2 ^Cl 2 +20 2 . 

Chloric acid is a powerful oxidizing agent. For example, it 
changes lead sulfide to lead sulfate (167). This operation is 
usually carried out by adding a few crystals of sodium chlorate 
and dilute hydrochloric acid to the black lead sulfide. The 
dark color is seen to change slowly to the white of the sulfate: 

3 PbS+4HC10 3 -> 3PbS0 4 + 4 HCl. 

355. Perchlorates and Perchloric Acid. — Perchlorates are 
formed by heating chlorates gently (306, 353). 

Sodium perchlorate, NaC10 4 , and potassium perchlorate, 
KC10 4 , are white crystalline salts. They decompose completely 
into chlorides and oxygen at dull-red heat. For example, 

NaC10 4 ->NaCl+20 2 . 



Oxidation and Reduction 213 

Ammonium perchlorate, NH 4 C10 4 , is made by neutralizing 
perchloric acid, HC0 4 , with ammonia. It is used as an oxidizing 
agent and as a very powerful explosive. 

When powdered sodium or potassium perchlorate is mixed 
with concentrated sulfuric acid and cautiously heated in a small 
retort (104, Fig. 24), perchloric acid, HC10 4 , is distilled from the 
mixture. This experiment should not be made by the student, 
as it might result in an explosion in unskilled hands. 

NaC10 4 +H 2 S0 4 ->NaHS0 4 +HC10 4 . 

Perchloric acid is a colorless liquid. It is a violent oxidizing 
agent, as shown by the fact that a drop of the acid will set fire to 
filter paper. The diluted acid is now coming into use in labo- 
ratories as an oxidizing agent, and also for the purpose of pre- 
cipitating potassium perchlorate in the quantitative analysis 
of potassium. 



CHAPTER XVI 



HEAT AND ENERGY 

356. Heat of Combustion. — Since coal, wood, and fuel gas 
are burned ordinarily in order to produce heat rather than as 
a means of obtaining their products of combustion, carbon 
dioxide and water, it becomes a matter of importance to discover 

how much heat is produced 
in the burning of a known 
weight of a given substance. 
The unit of heat is the 
calorie (in), which is the 
amount of heat required to 
raise the temperature of one 
gram of water one degree centi- 
grade. The amount of heat 
produced by the burning of 
one formula weight of a pure 
substance is called its heat 
of combustion. The heat of 
combustion of a solid is de- 
termined by burning a known 
weight of it within an appa- 
ratus of special design, called 
a bomb calorimeter. 

357. The Bomb Calorimeter. — This apparatus, illustrated in 
Fig. 44, consists of a heavy-walled metallic bomb with a gas- 
tight cover, surrounded by a vessel of water. The latter is con- 
tained in a larger vessel with walls of heat-insulating material. 
A weighed amount of substance whose heat of combustion is 
to be found is placed in the crucible of the bomb, which is filled 
with oxygen gas. The substance is then ignited by heat from 
a wire which carries an electric current. The temperature of 
the water surrounding the bomb is measured accurately before 











r 








_~ ~ 


1! 


k 


_J 


1 






_T\ 




^ 




fvi 


■-: 




I / JF 




V V V 







































Fig. 44 



214 



Heat and Energy 



215 



and after the burning, and the number of calories of heat pro- 
duced is calculated from the rise of temperature and the weight 
of water actually heated, plus the water equivalent of the bomb, 
etc. The water equivalent is the amount of water which has 
the same heat capacity as the bomb and other heated parts of 
the apparatus. Some typical results of measurements of heats 
of combustion are shown in Table XL The values are given 
to the nearest hundred, since this is about the limit of accuracy 
in such measurements. 

TABLE XI 



Substance 


Calories per gram 


Formula 


Heat of Combustion 


Carbon 


8,130 
34,400 

2,200 
11,900 

2,430 


C =12 g. 
H 2 = 2g. 

S =32 g. 

C 2 H 2 = 26g. 

CO =28g. 


97,600 
68,800 


Hydrogen 


Sulfur 


70,400 

315,400 

68,200 


Acetylene 


Carbon monoxide 







Since one formula weight of a gaseous substance has a volume 
22.4 liters, the heats of combustion of H 2 , C 2 H 2 , and CO are 
the amounts of heat produced in the burning of equal volumes 
of these gases. It will be seen that the heat of combustion of 
C 2 H 2 * is very large (nearly rive times that of hydrogen) . This 
accounts in part for the very high temperature of the oxy acety- 
lene flame (315). 

358. The British Thermal Unit, B.T.U.— In engineering 
practice quantities of heat are measured in British Thermal 
Units (B.T.U.) instead of in calories. This unit is the amount 
of heat required to raise the temperature of one pound of water one 
degree Fahrenheit. Since one pound equals 453 g., and i° F. = 5/9 
of i° C, it follows that 1 B.T.U. = 252 calories. The heat pro- 
duced in burning coal, coke, and fuel gas is called its calorific 
power. It is usually stated in terms of B.T.U. per pound of 
fuel. 

359. Composition and Calorific Power of Fuel. — Since the 
value of fuel is directly dependent on its calorific power, the 
testing of fuel is a matter of great practical importance. In 
testing coal it is customary to determine the moisture, volatile 



2l6 



Introduction to General Chemistry 



matter, "fixed carbon," and ash in addition to the calorific 
power. The "fixed carbon" is the non- volatile residue left 
when all volatile matter is driven off at a bright-red heat in the 
absence of air, less the ash contained therein. The calorific 
power is usually expressed in terms of B.T.U. per pound of fuel, 
or per cubic foot in the case of gases. Table XII gives some 
results for a variety of solid fuels. 



TABLE XII 



Kind of Fuel 



Lackawanna anthracite coal 

Pocahontas coal 

Indiana bituminous coal . 

Coke 

Lignite 

Oak wood 

Pine wood (resinous) .... 
Crude petroleum 



Percentage Composition 



Volatile 

Matter 



5 
18 

35 

0-5 
38 



Fixed 
Carbon 



74 
50 
QO 

Si 



Ash 



/ 
6 

9 
4 

0.4 
0.4 



Calorific Power 



Calories 
per gram 



7,724 
8,760 
8,080 
7,900 
7,200 
4,600 
5,000 
11,520 



B. T. U. 

per pound 



13,900 
15,680 
14,540 
14,200 
13,000 
8,300 
9,100 
20,736 



Table XIII gives the calorific power of some typical fuel 

gases. 

TABLE XIII 

Calorific Power in B.T.U. per Cubic Foot 

Kokomo, Indiana, natural gas 1 ,000 

Pittsburgh, Pennsylvania, natural gas 1,150 

Coal gas 650 

City of Chicago gas '. 600 



360. The Evaporation of Water and the Production of 
Steam. — We can easily calculate the amount of fuel theoretically 
needed to change water at ordinary temperature into steam. 
If one gram of water at 20 is heated to ioo°, 80 calories of heat 
are required, and in addition 540 calories are needed to change 
this hot water into steam. The total is 620 calories. Since 
the burning of one gram of coal produces about 8,000 calories, 
if all this heat were utilized it would be sufficient to evaporate 
(8,000-f- 620) 13 g. of water. In practice much heat is lost to 
the surroundings, as well as in the hot smoke which goes up the 



Heat and Energy 217 

smokestack. Engineers consider that it is good practice to 
evaporate 8 g. of water with 1 g. of coal. Therefore one pound 
of good coal will change 8 lb., or about 1 gal. (8.3 lb) of water 
at ordinary temperature into steam at ioo°. 

361. Heat of Reaction and Heat of Formation. — We have 
already frequently observed that numerous reactions other than 
combustions in oxygen (air) produce much heat. Among such 
are the reactions of chlorine with hydrogen (244), phosphorus 
(247), antimony (246), and turpentine (248); and water with 
sulfuric acid (93), potassium (106), and calcium oxide (150). 
The heat produced in these and other reactions may be measured 
in suitably constructed calorimeters and the results expressed 
most conveniently by stating the amount of heat given out in 
the reaction of formula weights of the uniting substances; or 
in the formation of one formula weight of the product. Thus 
the heat of reaction of CaO and H 2 may be written 

CaO+H 2 0->Ca(OH) 2 +5,ioo cal. 

and the heat of formation of water from its elements 

H 2 +|0 2 ->H 2 0+68,8oo cal. 

362. Heat of Neutralization.— The union of acids and bases 
to form salts and water always gives out heat. In fairly dilute 
solutions the amount of heat given out when one formula weight 
of water is so formed is almost exactly the same for many acids 
and bases. For example. 

HC1, NaOH = 13,700 cal. 
HC1, KOH =13,700 " 
HNO3, NaOH = 13,700 " 
HNO3, KOH =13,700 " 

This regularity is indeed striking and must mean close similarity 
in the processes of these reactions. How chemists interpret this 
phenomenon will be considered in chapter xviii. 

363. The Law of Constant Heat Summation. — Let us now 
consider the following question: If equal quantities of a given 
substance can be changed into the same product by two different 



218 Introduction to General Chemistry 

ways, will the amounts of heat produced be the same in the two 
cases? Carbon, for example, gives carbon dioxide when it is 
burned, 

C+0 2 ->C0 2 , 

but in a deficiency of oxygen the product is carbon monoxide, 

2C + 2 -»2CO. 

Carbon monoxide is a colorless gas which burns readily, giving 
carbon dioxide, 

2CO+0 2 ->2C0 2 . 

Therefore it is possible to change given weights of carbon and 
oxygen into carbon dioxide in two different ways. The heats 
of combustion are as follows: 

First Way 
C+|0 2 ^C0 +29,400 cal. 
CO+|0 2 ->C0 2 +68,2oo cal. 



Sum 97,600 cal. 

Second Way 
C+0 2 ->C0 2 +97,6oo cal. 

These results show that if 12 g. of carbon (C = i2) unite with 
32 g. of oxygen (0 2 = 32 and ^0 2 = 16) the total heat produced is 
the same no matter in which way the union occurs. 

Another illustration is found in the formation of a solution 
of ammonium chloride, NH 4 C1, from NH 3 and HC1 gases. 
This reaction can take place in two ways: 

First Way 
NH 3 (gas)+HCl (gas)->NH 4 Cl (solid) +42,000 cal. 
Heat absorbed in dissolving the NH 4 C1 in water = — 3,900 cal. 

Excess of heat produced over heat absorbed = 38,100 cal. 

Second Way 

Heat of solution of NH 3 in water = 8,400 cal. 

Heat of solution of HC1 in water = 17, 300 cal. 

Heat of neutralization of the two solutions = 12,400 cal. 



Total heat produced = 38,100 cal. 



Heat and Energy 219 

Innumerable cases like the two here given in illustration have 
led to the Law of Constant Heat Summation (Law of Hess). 
The heat produced or absorbed in the change of given substances 
into the same final products {in the same physical state) is the same, 
by whatever way the changes occur. 

That the heat of a given reaction is dependent on the physical 
state of the reacting substances and products is illustrated by 
the following example: 

CaO+H 2 (liquid)->Ca(OH) 2 (solid) + 15,100 cal. 
CaO+H 2 (ice) ->Ca(OH) 2 (solid) + 13,700 cal. 



Difference =1,400 cal. 

The difference, 1,400 cal., is due to the fact that it requires this 
amount of heat to change one formula weight of ice into water 
(18X79 = 1422) (118). 

364. Heat Produced in Slow Oxidation. Spontaneous Com- 
bustion. — Numerous experiments have proved that the amount 
of heat formed in a given reaction is just the same whether the 
change takes place slowly or rapidly. The decay of wood leads 
ultimately to the production of carbon dioxide and water, the 
same products as those formed when wood is burned. During 
the decay of wood, heat is produced so slowly that its formation 
is usually not perceptible by ordinary observation. Coal also, 
when exposed to the air, slowly oxidizes. In so doing it often 
loses an appreciable part of its heating value before it is burned. 
The depreciation on this account in the value of stored coal is a 
matter of considerable importance. 

If coal (especially bituminous coal) in small lumps and con- 
taining much dust is heaped in immense piles, such as are seen 
in coal yards, the heat produced by the slow oxidation does not 
escape readily from the bottom layers of the pile. The result 
is a gradual rise of temperature. At the higher temperature 
•oxidation and therefore heat production go on still faster, since 
usually enough air can diffuse in to keep up the supply of oxygen. 
Finally the temperature may rise so high that the pile of coal 
actually takes fire at the surface, where there is of course an 
unlimited supply of oxygen. Fire originating in this way is 



220 Introduction to General Chemistry 

said to be due to spontaneous combustion. The loss of coal 
through such fires was a very serious feature of the "coal famine" 
of 191 7-18. Some smoke is seen issuing from the majority of 
large piles of low-grade coal in the Chicago district, thus indi- 
cating more or less fire beneath. It is almost impossible to 
extinguish fire in a very large coal pile. The best way to prevent 
serious rise of temperature in coal piles is to provide numerous 
air shafts in the pile, by means of which warm air can escape. 
This does not entirely prevent oxidation but keeps the tempera- 
ture down to a point where the oxidation is not dangerously 
fast. 

It is a popularly known fact that " greasy" rags will often 
catch fire spontaneously. As a matter of fact such fires originate 
usually in rags soaked in oils used in paint or varnish, especially 
linseed oil or turpentine. The "drying" of paint and varnish 
is not a process of evaporation as much as one of oxidation of 
the oil used. These paint and varnish oils readily unite with 
oxygen to form solid products. In this process heat is produced. 
In a pile of rags, etc., covered with such oils sufficient rise of 
temperature may occur to cause spontaneous combustion. For 
this reason greasy rags, etc., should never be left where they 
can do damage if they take fire. 

365. Dust Explosions. — When the air is filled with the dust of 
coal, wood, flour, or other combustible substance a flame will 
often start a combustion which will spread with explosive 
rapidity. Appalling explosions have occurred from such causes 
in coal mines, wood- working factories, and flour mills. Even 
dust which is at rest in such places is blown into the air by the 
on-coming explosion wave and is thus changed to an explosive 
dust and air mixture. It is easy to see that a dust explosion is 
due to the extremely rapid burning of minute particles, each 
surrounded by an abundance of oxygen. Dust explosions are 
best prevented by keeping mines, mills, etc., free from accumula- 
tions of dust. 

366. Modes of Heat Production in Physical and Chemical 
Changes. — We have now learned that heat is produced (or 
absorbed) in a variety of physical and in all chemical changes. 



Heat and Energy 221 

The following seven modes of heat production (or absorption) 
have been studied: 

1. Latent heat of fusion (melting) (118). 

2. Latent heat of evaporation (115). 

3. Heat of solution (127). 

4. Heat of combustion (356). 

5. Heat of formation (361). 

6. Heat of reaction (361). 

7. Heat of neutralization (362). 

The first three modes have to do with physical changes of the 
sort known as changes of state ; the last four are due to chemical 
changes. All changes of state and many chemical changes are 
reversible processes. In every reversible process, if heat is 
given out when the change proceeds in one direction, heat is 
absorbed in equal amount when the change proceeds to an equal 
extent in the opposite direction. A change which results in the 
production of heat is called an exothermic change; one in which 
heat is absorbed is an endothermic change. 

367. Heat Production and Equilibrium. — In chapter xiii 
(288) the effect of temperature on equilibrium was discussed 
briefly. With respect to the change of solubility it was stated 
that raising the temperature causes that change of solubility to 
occur which involves an absorption of heat. We also saw (288) 
that for chemical equilibrium raising the temperature causes 
the equilibrium to shift in the direction that involves an absorption 
of heat. These laws are entirely general and apply to all 
changes of state and all chemical changes. 

In the shift of equiHbrium which occurs with change of 
temperature the fraction of the reacting substances transformed 
to new products is determined, in a given case, by the change 
of temperature (measured in degrees). The amount of heat 
(in calories) absorbed (if the temperature is raised) or given out 
(if the temperature is lowered) is determined by the amount of 
material transformed. An example will make the matter clearer. 

Hydrogen and iodine vapor react partially in the neighbor- 
hood of 400 to give hydrogen iodide (264, 281, 288): 

H 2 +I 2 ± 5 2HI+i,ooo cal. 



222 Introduction to General Chemistry 

i 

This equation means that the formation of two formula weights 
of HI from H 2 and I 2 (vapor) at about 400 takes place with 
the liberation of 1,000 cal. of heat or 500 cal. for each formula 
weight of HI produced. The following table shows the propor- 
tions of molecules in the equilibrium mixture at 370 and 440 : 





H, 


la 


HI 


Total 


370 


IOO 


IOO 


800 


IjOOO 


440 


no 


no 


780 


1,000 



We see that if the temperature is raised from 370 to 440 , 20 
molecules of HI out of a total of 1,000 molecules (2 per cent of 
the whole) change into H 2 and I 2 . If the total amount of 
material in the mixture is that resulting from one formula weight 
each of H 2 and I 2 (equivalent to two formula weights of HI), 
and if 2 per cent of the whole number of molecules change into 
H 2 and I 2 , the heat absorbed is 0.02 X 1,000 cal. = 20 cal. 

368. Work and Energy. — The terms work and energy have 
very definite meanings in science. The subject of physics is 
largely concerned with these very important matters ; and since 
it is assumed that the student has already studied physics, an 
elementary discussion of these very important topics is unneces- 
sary. We may, however, briefly summarize some of the more 
prominent points. The typical example of work in the physical 
sense is the lifting of a weight. The scientific unit of work is 
the gram centimeter, which is the work required to lift one gram 
one centimeter. The amount of work done in lifting a weight is 
the product of the force required (which in this case is equal 
to the weight in grams) and the vertical distance measured in 
centimeters. Thus the lifting of 600 g. to a height of 30 cm. 
requires the doing of 600X30 = 18,000 g. cm. of work. The 
weight of 600 g., having been lifted 30 cm., could do work to 
the extent of 18,000 g.cm. in descending 30 cm. It is said to 
have the power to do this amount of work. Now power to do 
work is called energy, and therefore it has 18,000 g.cm. of 
energy. Two kinds of energy are recognized: potential energy, 
as possessed by a weight which may do work on descending, and 
kinetic energy, or the energy of a body in motion. It requires 



Heat and Energy 



223 



work to set a body in motion, and conversely a body in motion 
is able to do work. 

369. The Mechanical Equivalent of Heat. — Heat is also a 
form of energy, because heat is able to do work. A steam engine 
is merely a machine which converts the heat of burning coal into 
kinetic energy. The change of kinetic energy into heat may be 
observed on every hand: anything that restrains or stops the 
motion of a moving body converts part or all of its kinetic energy 
into heat. We measure energy in gram centimeters and heat 
in calories, and if heat is a form of energy then the calorie, like 
the gram centimeter, must be an energy unit. It will at once 




w 



Fig. 45 



be asked: Do these units represent equal amounts of energy? 
In other words, will one gram centimeter of work produce one 
calorie of heat? If not, how many gram centimeters are required 
to produce one calorie? This question was first answered by 
Joule in 1840. 

370. Joule's Experiment. — In Joule's experiment, with 
apparatus shown in Fig. 45, a weight, W, attached to a cord 
wound on a cylinder, in slowly descending turns a stirrer which 
is surrounded by water in a calorimeter, C. The water, which 
restrains the motion of the stirrer, becomes warmer, owing to 
the change of work into heat. The amount of work in gram 
centimeters done in heating the water is the product of the mass 
in grams of the weight and the distance of its descent in 



224 Introduction to General Chemistry 

centimeters. The amount in calories of heat produced is the 
product of the rise in temperature in degrees C. and the mass, 
in grams, of water plus the water equivalent of the heated parts 
of the calorimeter. 

By means of this apparatus Joule found pretty closely the 
number of gram centimeters of work equivalent to one calorie 
of heat. More refined work since then has shown that one 
calorie is equal to 42,700 g.cni. This ratio is called the mechanical 
equivalent of heat. This means, for example, that one gram 
falling 42,700 cm. (a little over a quarter of a mile) produces one 
calorie. 

371. The Conservation of Energy. — At the time Joule began 
his experiments in 1840 it was not at all clear that the amount 
of heat produced by a given amount of work (kinetic or potential 
energy) was definite. It seemed possible, if not probable, that 
different modes of changing work into heat would give different 
values for the mechanical equivalent. So Joule used not only 
the method and apparatus already described but also two others. 
His three methods and the mechanical equivalent of one calorie 
were as follows: (1) stirring water in a brass vessel with a brass 
paddle, 42,400 g.cm.; (2) stirring mercury in an iron vessel 
with an iron paddle, 42,500 g.cm.; (3) rubbing two iron rings 
together under mercury, 42,500 g.cm. 

The very close agreement of the results of the three experi- 
ments led Joule to conclude that the amount of heat produced 
by a given amount of work is always the same, by whatever way the 
work is changed into heat. This result has been amply confirmed 
by all later experiments and experience. When work of any 
kind (mechanical energy, either kinetic or potential) is changed 
into heat there is no real loss or destruction of energy, since the 
heat produced is also energy in another form and exactly equal 
in amount to the work done in producing it. This conclusion 
is concisely stated in the Law of the Conservation of Energy: 
Energy is indestructible. 

Just as the law of the conservation (indestructibility) of 
matter (21) is the foundation stone of the science of chemistry, 
so, similarly, this law of the conservation (indestructibility) 



Heat and Energy 225 

of energy is the solid rock upon which the whole structure of the 
science of physics rests. 

372. Other Forms of Energy. — We have defined the term 
energy as the power of doing work; and since heat is also a form 
of energy, we might extend the definition so as to read: Energy 
is the power to do work or produce heat. According to this defini- 
tion of energy it is obvious that light and even sound and espe- 
cially electric currents are also forms of energy, since each of 
these by appropriate means can produce work or heat. 

373. Chemical Energy. — For the chemist an important ques- 
tion now arises : What shall be said of the source of energy that 
produces the great heat of a burning substance? This question 
is somewhat like the one, What is the source of energy of a 
" wound-up" watch spring? To wind up the spring a certain 
amount of work must be done. Is it not reasonable to say that 
the energy used in winding up the spring has been "stored up" 
in the coiled spring? If so, we may say that this energy is 
changed into potential energy, just as we say that the energy 
required to lift a weight is changed into potential energy and 
can be regained as useful work then the weight is allowed to 
descend. Reasoning somewhat similarly, we may conclude that 
the energy given out as heat in the burning of hydrogen, for 

which we have 

H 2 +J0 2 ->H 2 0+68,8oo cal, 

comes from some form of potential energy which has been stored 
up in the two gases. This conclusion is rendered highly probable 
by reason of the fact that by means of an electric current (elec- 
trical energy) we can decompose water into hydrogen and oxygen. 
Since the electrical energy disappears and very little heat is 
formed, we may very reasonably conclude that it has been 
changed into some sort of potential energy stored up in the two 
gases formed from the water. The form of potential energy 
stored up in chemical substances and liberated when they react 
is called chemical energy. 

374. The Sun as a Source of Energy. — It will be interesting 
to trace some familiar form of energy through various trans- 
formations back to its source. Take, as an example, the energy 



226 Introduction to General Chemistry 

given out as light and heat by an electric lamp. The energy 
comes to the lamp as an electric current having electrical energy. 
This electrical energy was produced in a dynamo or generator, 
the armature (the moving part) of which was turned by a steam 
engine. The kinetic energy of the engine was derived from hot, 
compressed steam produced from water by the burning of coal 
which has resulted from the slow transformation of vegetable 
matter. 

Plants derive nearly all of their substance from water and 
the carbon dioxide of the air under the influence of the light and 
heat of the sun. A great deal of energy is taken up by plants 
as light and heat and is stored as chemical energy in the sub- 
stances composing them, as well as in the oxygen which is set 
free by the growing plant. Recapitulating, we see that the 
light and heat from the sun are changed by growing plants into 
chemical energy; this energy is largely conserved when plants 
are changed into coal. When the coal burns, its chemical 
energy, supplemented by that of the oxygen of the air, is changed 
into heat, which is in turn changed into kinetic energy in the 
steam engine. .The kinetic energy of the engine is then changed 
by a dynamo into electrical energy, and the latter produces in 
the lamp heat and light. 



CHAPTER XVII 
THE IONIC HYPOTHESIS 

375. The Ionic Hypothesis. — This chapter will treat of the 
properties and behavior of acids, bases, and salts and aims to 
show how a supposition called the ionic hypothesis furnishes a 
satisfactory explanation of many facts. 

376. The Two Parts of a Salt. — It must have been noticed 
that a salt is made up of two parts, the metallic or basic part and 
the non-metallic or acidic part. The latter may be an element 
like chlorine in sodium chloride; or it may be a radical (147) 
like S0 4 , which is contained in every sulfate. The name of a 
salt always indicates the parts of which it may be considered as 
being made up. Thus potassium nitrate, KN0 3 , is composed 
of potassium and nitrate radical, N0 3 ; and calcium carbonate, 
CaC0 3 , of calcium and carbonate radical, C0 3 . 

377. The Two Parts of an Acid. — Every acid may also be 
considered as made up of two parts, one of which is hydrogen 
and the other the characteristic acid radical of that acid. For 
example, sulfuric acid, H 2 S0 4 , may be considered to consist of 
hydrogen and sulfate radical, S0 4 ; and phosphoric acid, H 3 P0 4 , 
to consist of hydrogen and phosphate radical, P0 4 . For this 
reason S0 4 ' and P0 4 may be called acidic radicals. Dilute 
solutions of pronounced acids have a sour taste. Since hydro- 
gen is the only constituent which all acids have in common, we 
may reasonably conclude that the sour taste is due to the H 
radical. 

378. The Two Parts of a Base. — A base is usually the 
hydroxide of a metallic element, and it may therefore be con- 
sidered as made up of two parts, the metal and the hydroxyl 
radical, OH. Thus sodium hydroxide, NaOH, consists of 
sodium and hydroxyl radical, OH. Ammonium hydroxide, 
NH 4 OH, may be considered as made up of ammonium radical, 
NH 4 , and hydroxyl. Consequently the ammonium radical may 

227 



228 Introduction to General Chemistry 

be called a basic radical. It is the only basic radical that we 
have studied, although many others are known. 

379. The Process of Neutralization. — The two following 
equations represent typical cases of neutralization: 

NaOH+HCl->NaCl+H 2 ; 
NH 4 OH+HN0 3 ->NH 4 N0 3 +H 2 0. 

We notice that in each case the salt which is formed is made up 
of two parts, one of which comes from the base, the other from 
the acid. In each case water, whose formula may be written 
HOH, is also formed. We might call water hydrogen hydroxide 
and think of it as being made up of two parts hydrogen and 
hydroxyl radical. The process of neutralization consists, there- 
fore, merely of an exchange of partners, so to speak, on the part 
of the base and the acid. 

As a matter of fact, not only can neutralization be represented 
in this way, but most reactions in water solution between acids, 
bases, and salts which do not involve oxidation or reduction 
may be regarded as an exchange of the partners of the reagents 
initially used. This will be made clear by the following examples. 

380. Reactions of Barium Salts with Sulfates. — If we add 
dilute sulfuric acid to a dilute solution of barium chloride a white 
precipitate of barium sulfate is formed, 

H 2 S0 4 +BaCl 2 ->BaS0 4 +2HCL 

A precipitate of barium sulfate also results when a solution 
of any barium salt is added to a solution of any sulfate, as 
illustrated by the following equations: 

K 2 S0 4 +Ba(N0 3 ) 2 ^BaS0 4 +2KN0 3 , 
CuS0 4 +BaBr 2 -^BaS0 4 +CuBr 2 . 

This is so generally true that the formation of a precipitate 
of barium sulfate upon the addition of a solution of a barium 
salt to some other solution shows that this second solution con- 
tains the sulfate radical, S0 4 , in the form either of a sulfate or 
of sulfuric acid. We say therefore that the formation of a 
precipitate of barium sulfate when a solution of a barium salt is 
added to another solution is a test for the sulfate radical. It is 



The Ionic Hypothesis , 229 

important to note that it is the S0 4 radical, and not sulfur or 
oxygen or a combination of the two in some other proportion, 
that responds to this test. A solution of hydrogen sulfide, 
H 2 S (339), which may be considered as being made up of two 
parts, hydrogen and sulfur, does not give a precipitate of any 
sort with a solution of a barium salt. Furthermore, pure dilute 
sulfurous acid, H 2 S0 3 (340), which is made up of hydrogen and 
sulfite radical, S0 3 , does not give a precipitate when mixed with 
a barium salt solution. 

381. Reactions of Simple Lead Salts. — Lead sulfate, PbS0 4 , 
is also a white insoluble salt. If we add a solution of any 
simple lead salt to a dilute solution of sulfuric acid or any soluble 
sulfate, we obtain a white precipitate of lead sulfate, 

Na 2 S0 4 +Pb(N0 3 ) 2 ->PbS0 4 +2NaN0 3 . 

In this case, just as in the precipitation of barium sulfate, it is 
the sulfate radical, S0 4 , which has united with the lead to form 
the precipitate. 

It is also of equal interest to note that if the nitrate of barium 
or of lead is used, the nitrate radical, N0 3 , is left in combination 
with the basic element or radical which was originally combined 
with the sulfate radical. 

382. The Reaction of Silver Salts with Chlorides. — We have 
already learned that a solution of silver nitrate reacts with 
hydrochloric acid or a chloride to give a white precipitate of 
silver chloride : 

AgN0 3 +NaCl->AgCl+NaN0 3 . (169) 

A solution of any simple silver salt reacts similarly with hydro- 
chloric acid or any chloride, so that we may think of the reaction 
as characteristic and call it a test for silver salts. This reaction 
is specifically that of the chloride radical; for if we add silver 
nitrate solution to a solution of potassium chlorate, KC10 3 
(353)? no apparent change is observed; certainly no silver 
chloride is formed, otherwise the latter, being insoluble, would 
separate out as a white precipitate. This shows that chlor- 
ine in the chlorate radical, C10 3 , behaves entirely differently 



230 ► Introduction to General Chemistry 

from chlorine in the form of a chloride. We also find that 
solutions of perchlorates, of which potassium perchlorate, 
KCIO4 (355), is an example, do not give precipitates with solu- 
tions of silver salts. It is possible to make both silver chlorate, 
AgC10 3 , and silver perchlorate, AgC10 4 , by methods which we 
need not consider at present, and it is found that these salts are 
entirely different from silver chloride, and that both are easily 
soluble in water. 

This brief review of reactions, most of w T hich have already 
been studied, is sufficient to illustrate the subject in hand, but 
many other examples of the same principle will be found in the 
previous chapters. 

383. Summary and Conclusions. — The observations which 
we have made are typical for all acids, bases, and salts. Each 
may be shown to be made up of two parts. In the examples we 
have studied these are, on the one hand, hydrogen, a metal, or 
the ammonium radical, and, on the other, hydroxyl, a halogen, 
sulfur, or an acid radical. Hydrogen is one of the two parts of 
every acid, and hydroxyl one of the two parts of every base. 
In chemical reactions between acids and bases, acids and salts, 
bases and salts, and between two salts (where oxidation and 
reduction do not occur) the two substances simply exchange 
parts. This kind of chemical change is called double decompo- 
sition (337) or metathesis. The chemical reactions which acids, 
bases, and salts give are in reality only the reactions of their parts. 

Finally it should be noted that the recombination of these 
parts always takes place between the basic or metallic part on 
the one hand and the acidic or the non-metallic part on the 
other. Double decompositions in water solutions never give 
compounds such as KNa or C1S0 4 . This is a striking observa- 
tion, and the fact that we have as yet no explanation for it 
warns us at once that we must go farther in our observations 
to understand even the most commonplace of these reactions. 

384. Double Decomposition and Electrical Conductivity. — 
Along with the ability to undergo double decompositions, acids, 
bases, and salts in their water solutions have the property of 
conducting the electric current. If we set up a battery, a 



The Ionic Hypothesis 



231 



galvanometer, and a salt solution in the manner shown in 
Fig. 46, using platinum electrodes and a sufficient number of 
dry cells or other source of current to give a suitable deflection 
of the galvanometer, we shall find that if we replace the solution 
by distilled water practically no current will be indicated by the 
galvanometer. We also find that if dry salt is substituted for 
the solution no current will pass. If now we pour distilled 
water on the salt while the latter is still in contact with the 
electrodes, a current begins to pass through the solution of salt 
which is quickly formed. 



if 



£4 



&£, 



Fig. 46 

These results lead to the conclusion that neither pure water 
nor dry salt conducts the current appreciably compared with 
the solution formed from salt and water. All other soluble 
salts behave similarly. It is also easily discovered by experi- 
ment that dry (water-free) bases and acids are no better con- 
ductors than dry salts, although water solutions of acids and 
bases are good conductors. Water solutions of other substances 
than acids, bases, and salts do not conduct electricity any better 
than does pure water. 

The close connection which always exists between electrical 
conductivity and the ability of a mixture to undergo double 
decomposition is illustrated by the following experiment. Ferric 
sulfate and sodium carbonate can be mixed dry without any 
apparent change; but let the mixture once be wet with water, 
immediately a violent evolution of gas follows and a red 
precipitate of ferric hydroxide appears. That the mixture of 
the dry substances is a non-conductor is shown by placing it 
in the dry beaker. Fig. 46. No current passes, but when water 
is added the substances dissolve, and the solution so formed 
conducts the current. At the same time the chemical reaction 



232 Introduction to General Chemistry 

begins vigorously. Since chemical reactivity and electrical con- 
ductivity seem therefore to go hand in hand, we shall next study 
the behavior of solutions of acids, bases, and salts when an electric 
current is passed through them. 

385. The Electrolysis of Solutions. — We have already learned 
that the electrolysis of concentrated hydrochloric acid sets free 
hydrogen and chlorine (43), and that the electrolysis of common 
salt (238) yields these same gases and in addition forms sodium 
hydroxide. In the case of hydrochloric acid, electrolysis simply 
causes the separation of the two constituents, 

HC1->H+C1. 

On being set free the atoms of the two elements each form 
diatomic molecules, thus, 

2H->H 2 , and 2C1-»C1 2 . 

These last reactions are doubtless secondary; and for the sake 
of brevity, in the examples of electrolysis that follow, reactions 
of this kind will be indicated by separate equations without 
further comment. 

In the case of the electrolysis of common salt it seems pos- 
sible, as already explained (238), that the first change is a 
separation into sodium and chlorine, thus : 

NaCl->Na+Cl, 
2CI-XX 

The sodium then reacts with the water present to form sodium 
hydroxide and hydrogen: 

Na+H 2 0->NaOH+H, 
2H->H 2 . 

Whether this is the only possible explanation of the way the 
changes take place can best be discussed later; but it can be 
pointed out here that the foregoing equation would demand that 
the sodium hydroxide should be formed at the electrode at which 
the hydrogen is given off; and this is, in fact, the case. When 
hydrogen is set free in the electrolysis of any substance it always 



The Ionic Hypothesis 233 

appears at the negative electrode or cathode, while chlorine is 
liberated only at the positive electrode or anode. 

386. The Electrolysis of Copper Salts. — If a solution of 
cupric chloride, CuCl 2 , is electrolyzed between platinum poles 
or electrodes, copper is deposited on the negative pole and 
chlorine gas is set free at the positive pole. Here again, as in 
the case of hydrochloric acid, we have the simplest possible 
kind of a change, as represented by the following equation: 

CuCtf»Cu+2Cl, 
2C1->C1 2 . 

If copper sulfate, CuS0 4 , is electrolyzed, a plating of metallic 
copper is again formed on the negative electrode, while from the 
positive electrode oxygen gas is given off. Examination of 
the products after electrolysis shows that sulfuric acid is con- 
tained in the solution surrounding the positive electrode. In 
fact, if the electrolysis is continued until all the copper in the 
original solution is deposited, all the sulfate radical of the 
original copper sulfate is changed into sulfuric acid, and this 
acid is contained in the part of the solution surrounding the 
positive electrode. The formation of sulfuric acid and oxygen 
may be explained by supposing the copper sulfate to be separated 
by the electric current into copper and sulfate radical, S0 4 , and 
that the latter reacts with water to form sulfuric acid and oxygen: 

S0 4 +H 2 0->H 2 S0 4 +0, 
20->0 2 . 

387. The Electrolysis of Silver Nitrate. — If an electric cur- 
rent is passed through a solution of silver nitrate, AgN0 3 , silver is 
deposited on the negative electrode and oxygen and nitric acid 
appear at the positive electrode. Probably silver nitrate is 
first separated into silver and nitrate radical, N0 3 ; the latter 
then reacts with water to form nitric acid and oxygen : 

2N0 3 +H 2 0->2HN0 3 +0, 

20->0 2 . 

388. Summary. — In Table XIV we have summarized the" 
results just discussed, leaving out of consideration the secondary 



234 



Introduction to General Chemistry 



changes that often take place between the substance set free and 
the water. We see that the parts into which a substance is separated 
by electrolysis are the same as those which change partners in double 
decomposition reactions. 

TABLE XIV 
Immediate Products of Electrolysis 



Liberated at 


Original 


Liberated at 


Negative Plate 


Substance 


Positive Plate 


H 


HC1 


CI 


Na 


NaCl 


CI 


Cu 


CuCl 2 


2C1 


Cu 


CuS0 4 


so 4 


Ag 


AgN0 3 


N0 3 



389. The Terms Used in Electrolysis. — The phenomena of 
electrolysis were very carefully studied about 1834 by Faraday, 
who, as we shall see, discovered many important facts. It was 
Faraday also who invented the terms electrolysis, electrolyze, 
electrode, anode, and cathode. He called the solution the 
electrolyte, but at present we use this term to mean the dissolved 
substance. That part of the electrolyte which during electrolysis 
is deposited or set free at the anode or positive efectrode he 
called the anion. The other part, which goes to the cathode, he 
called the cathion. Frequently he had occasion to speak of 
anions and cathions together, and then he referred to them as 
the ions of the electrolyte. For example, the ions of copper 
sulfate are said to be copper and sulfate radical, because in 
electrolysis copper passes to and is deposited on one electrode, 
while the sulfate radical goes to the other. Of course at some 
time or other the radicals or the partners of the original elec- 
trolyte must have broken apart, otherwise they could not have 
arrived at poles distant from each other. 

390. Hydrogen and Metals as Cathions. — We may next con- 
sider how the composition of the ions of a substance can be 
discovered. 

In the case of such a simple substance as HC1 it is obvious 
•that the ions are hydrogen and chlorine, hydrogen being the 
cathion and chlorine the anion. Since all acids upon electrolysis 



The Ionic Hypothesis 235 

give off hydrogen at the cathode, we may conclude that hydrogen 
is the cathion of all acids. 

When salts are electrolyzed the metal is either deposited in 
metallic form on the cathode, as in the case of copper and silver 
salts, or it collects in the solution about the cathode in the form 
of hydroxide, as when common salt is used. These facts lead 
to the conclusion that the basic or metallic elements of salts are 
cathions. 

391. Acid Radicals as Anions. — On the other hand, the acid 
elements or radicals of acids and salts accumulate at the anode 
and are either given off as free elements, as in the case of chlorine, 
bromine, and iodine, or they react with water to form acids and 
oxygen, as in the case of sulfate and nitrate radicals. 

392. Ions and Chemical Reactions. — It would thus appear, 
from what has just been stated, that the ions of an acid or salt 
are the same as the two parts of which its chemical reactions 
show it to be composed. It may be added that there is good 
reason for thinking that the same statement also applies to 
bases. The cathion of a base is usually a metal; the anion 
is the hydroxyl radical. 

393. Positive and Negative Ions. — The cathode is the electro- 
negative electrode; to it go the cathions. Since it is well known 
that a negatively charged body repels another negatively 
charged body and attracts one which is positively charged, it is 
not unreasonable to attribute the movement of ions to electrical 
attraction, and to conclude that cathions are electropositively 
charged. Since the anode is electropositive, the anions are 
thought to be charged electronegatively. It is customary to call 
cathions positive ions, and anions negative ions. 

394. The Cause of the Union of Ions. — Attention has been 
called to the fact that in reactions in solution basic or metallic 
radicals unite only with acidic or non-metallic radicals (383), 
and that unions of basic radicals with one another never occur; 
that is, double decompositions in water solutions never give 
compounds such as KNa and C1S0 4 . We are now in position to 
explain these important facts. We have learned that the radicals 
of acids, bases, and salts are identical with their ions, and that 



236 Introduction to General Chemistry 

the ions are probably electrically charged, the basic or metallic 
ones being positively, the acidic or non-metallic negatively, 
charged. We can therefore summarize by stating that in 
reactions only ions of unlike electric charges unite with one another. 
The reason for this is that ions with unlike electric charges 
attract each other, and that those with like charges repel each 
other. The chemical union of ions is an electrical phenomenon 
and is due to the attractive force of unlike electric charges carried 
by the ions. 

395. Colors of Ions. — The student has doubtless already 
observed that, although most salts and their solutions are color- 
less, a considerable number are colored. A little investigation 
will show that very dilute solutions of equal concentration of 
all cupric salts of colorless acids are of the same shade and 
intensity of blue color. This fact leads us to believe that the 
blue color is due to copper ions, which is the only substance which 
all the solutions have in common. Moreover, we find that the 
colors of all dilute solutions of colored acids, bases, and salts 
can be ascribed to the colors of their ions. 

If the dilute solution of any acid, base, or salt is colorless, 
like pure water, we may conclude that its positive and negative 
ions are both colorless. If "a dilute colored solution of an elec- 
trolyte has one colorless ion we conclude that the observed 
color is that of the other ion. Thus we find that all dilute 
ferrous solutions (173, 331) are pale green and conclude that 
ferrous ion is pale green. Manganous salts (342) (for example, 
MnCl 2 and MnS0 4 ) give pale pink solutions, therefore positive 
Mn ion is pale pink. On the other hand, dilute solutions of all 
permanganates (343) are, like KMn0 4 , deep purple, and we 
conclude that negative Mn0 4 ion is purple. Similar reasoning 
leads us to conclude that negative Cr0 4 ion is yellow (345) and 
negative Cr 2 7 ion is orange (345), while positive Cr ion is violet 
(344). The color of a dilute solution is usually an indication of 
the nature of one of its ions. 

396. Colors of Molecules. Dissociation of Molecules into 
Free Ions. — Although dilute solutions of all cupric salts are 
blue the solid salts and also their concentrated solutions are in 



The Ionic Hypothesis 237 

several cases of a different color. Thus cupric chloride, CuCl 2 , 
in solid form and in concentrated solution is green, and cupric 
bromide, CuBr 2 , which forms almost black crystals, gives a 
concentrated solution which is dark brown; but if this brown 
solution is sufficiently diluted the color gradually changes to 
blue, finally reaching the same shade of color as that of any 
other equally dilute cupric solution. A simple explanation 
of these color changes is found in the assumption that the dark 
brown color is that of the molecules, CuBr 2 , while the blue color 
is due to Cu ions. From the fact that many dilute solutions 
of bromides are colorless we conclude that Br ions are colorless. 
By following up this idea we are led to a very remarkable con- 
clusion, namely, that molecules of CuBr 2 exist only in the solid 
state and in concentrated solutions but not to an appreciable 
extent in very dilute solutions. This is accounted for if we 
assume as the concentrated solution is diluted molecules gradually 
split up or dissociate into ions, thus : 

CuBr 2 ->Cu+2Br, 

so that in a dilute solution the substance exists largely as free Cu 
and Br ions. If we evaporate the dilute blue solution we again 
obtain a brown concentrated solution and finally brown crystals. 
We must therefore assume that the change is a reversible one, the 
ions reuniting to form molecules as the solution is evaporated. 
Further evidence of the existence of free ions is afforded by the 
experiments next to be considered. 



sa 




Fig. 47 

397. The Migration of Ions. — Let us take advantage of the 
color of ions to discover their behavior during the process of 
electrolysis. In the U-tube, Fig. 47, we may put a solution of 
a colored electrolyte in the lower layer and colorless electrolytes 



238 



Introduction to General Chemistry 



in the layers next to the electrodes. As colored electrolytes we 
may use copper nitrate or potassium permanganate. When the 
current is turned on, the boundary of each colored electrolyte 
slowly moves up into one of the colorless layers above it, just 
as we would expect if the colored materal is the free ion which 
carries a charge of electricity. Thus positive copper ion migrates 
toward the negative electrode, and negative permanganate ion 
migrates toward the positive electrode. We can carry out an 
experiment with a mixture of these two colored salts in the lower 
layer. The purple layer now shows on the side of the positive 
electrode, and the blue layer shows on the side of the negative 
electrode just as before. Thus we find that each ion migrates 
just as though the other were not there; and this, in fact, is 
just what we should expect if a dilute solution contains free 
ions formed by dissociation of its molecules. 

398. The Mechanism of Electrolysis. — We can now make a 
fairly complete picture of the mechanism of the- conduction of 
the current through a solution and of the accompanying elec- 
trolysis. We shall assume that in the dilute solution the dis- 
solved substance is partially dissociated into positive and 

negative ions. Fig. 48 
represents diagrammati- 
cally such a solution 
in which the two elec- 
trodes are dipped, con- 
nected with a pair of 
dry cells. The cells 
charge the electrodes, 
one positively, the other negatively. The influence of these 
charges is felt by the ions in the solution. The positive ions 
are attracted by the negative electrode and repelled by the 
positive electrode and in consequence migrate toward the 
former. The negative ions move in the opposite direction for 
similar reasons. 

When ions come into contact with the electrodes of unlike 
sign they give up their charges to the electrodes. This tends 
to discharge the latter, but the battery keeps them charged by 



- fee « (=> 
eeg q 
^©9 e© © 

© e e©e' 




Fig. 



The Ionic Hypothesis 



2 39 



continuously sending a current of electricity through the wires. 
A more detailed description of the mechanism of electrolysis, 
in terms of the newer views of the nature of electricity, will be 
given in chapter xx. 

399. Faraday's Laws of Electrolysis. — As the result of care- 
ful experimental investigation of the quantities of substances 
liberated during electrolysis, Faraday arrived at the following 
conclusions: 




«--^L 



r 



Fig. 49 

1 . The amount of a given substance, say hydrogen, set free by 
electrolysis is directly proportional to the quantity of electricity 
which is passed through the solution. 

2. The amount of a substance, hydrogen for example, which 
is liberated by a fixed quantity of electricity is the same, whatever 
the nature of the solution electrolyzed, provided that this substance 
and nothing else is liberated at the given electrode. These two 
statements are known as Faraday's Laws of Electrolysis. 

400. Two Electrical Units. — To understand these laws fully 
we must review briefly some fundamental facts about the elec- 
trical current so that we can appreciate what is meant by 
quantity of electricity. In the first place we know that if a 
current passes through a wire there is produced around the wire 
a magnetic field. If we attach a thread to the middle of a 
magnetized steel needle and suspend the latter above and 
parallel to a wire, then as soon as we cause a current of elec- 
tricity to pass through the wire the needle sets itself at an angle 

The greater the angle between the 



to the former, Fig. 49. 



240 



Introduction to General Chemistry 




Fig. 



needle and the wire the stronger is the magnetic field, and the 
stronger the current is said to be. It is on this principle that 
instruments are built to measure current strength. Of course, 
to measure anything we must first adopt some fundamental 
unit by comparison with which the measurement can be made. 
This was done in the case of the electric current on the basis of 
the strength of the magnetic field about a conductor, and this 
unit was called the ampere. The ammeter 
(Fig. 50) allows us to read, from the posi- 
tion occupied by the needle on its scale, just 
how many amperes of current are passing. 
The amperage tells us the strength of 
the current, but we must also know the 
time during which the current is allowed to 
pass if we are to know the amount of elec- 
tricity delivered at the terminals of the con- 
ductor, say at two electrodes. // a current of one ampere is 
allowed to flow one second it is said to deliver a unit quantity of 
electricity, and this unit is called the coulomb. 

401. Illustration of Faraday's Laws. — -The following facts 
will serve to illustrate the meaning of Faraday's laws. By the 
electrolysis of dilute acids hydrogen is set free at the negative 
electrode. In all cases the passage of 96,500 coulombs of 
electricity is required for the liberation of one gram of hydrogen. 
Since a current of one ampere delivers one coulomb per second, 
96,500 coulombs will be given by a current of one ampere in 
96,500 seconds, or 26.8 hours. A current of two amperes for 
the same length of time will liberate 2 g. of hydrogen, or one 
gram molecular weight of hydrogen (H 2 ), which as we know has 
a volume of 22.4 liters at o° and 76 cm. 

402. Discussion. — It is not surprising that if a one-ampere 
current will liberate 1 g. of hydrogen in 26 . 8 hours, a two-ampere 
current will liberate 2 g. of hydrogen in the same time, for this 
is the type of regularity which we have become accustomed to 
expect in nature. It is surprising, however, that the same 
amount of hydrogen is liberated by the same amount of elec- 
tricity from a solution of any dilute acid, and the fact that this 



The Ionic Hypothesis 



241 



is so must reflect some regularity in the phenomena of elec- 
trolysis, the cause of which we have still to discover. 

403. Faraday's Laws of Electrochemical Equivalents. — Let 
us now turn to cases of the liberation by electrolysis of elements 
other than hydrogen. Very careful experimentation has shown 
that by the passage of 96,500 coulombs of electricity through 
various solutions certain weights of elements are set free. These 
are given in Table XV. This table shows a most striking regu- 

TABLE XV 

Electochemical Equivalents 



Element 


Gram Atomic 
Weight 


Valence 


Weight Liberated 
in Grams 


Gram Atomic 
Weight -s- Valence 


Hydrogen 

Silver 


1 
108 

64 

66 
208 

56 

27 

35-5 

80 

16 


1 

1 . 

2 

2 

2 

2 

3 
I 

I 
2 


I 

108 

32 

33 

104 

28 

9 

35-5 

80 

8 


1 
108 


Copper 


32 

33 

104 


Zinc 


Lead 


Iron 


28 


Aluminum 

Chlorine 

Bromine 

Oxvgen 


9 
35-5 

80 
8 







larity: The weight of an element liberated in electrolysis by the 
passage of 96,500 coulombs of electricity is equal to the gram atomic 
weight of that element divided by its valence (col. 5). This weight 
is called the electrochemical equivalent of a given element or, 
more briefly, its equivalent weight. The electrochemical equiva- 
lents of the various elements are seen to be proportional to the 
weights of these elements which unite chemically with one 
another when union is possible; for example, 1 g. of hydrogen, 
104 g. of lead, or 9 g. of aluminum unite with 35 . 5 g. of chlorine, 
or 8 g. of oxygen. The discovery of facts such as those given 
in the table was made by Faraday, who stated his conclusion as 
the Law of Electrochemical Equivalents : The amounts of different 
substances liberated by the same quantity of electricity arc propor- 
tional to their equivalent weights. 

404. The Electric Charges of Ions. — The facts covered by 
Faraday's laws allow us to draw some interesting and significant 



242 Introduction to General Chemistry 

conclusions regarding the quantities of electricity composing 
the charges on single ions. If the 96,500 coulombs of electricity 
supplied at the negative electrode to release one gram of hydro- 
gen ion are used simply to neutralize the charge on one gram of 
that ion, we may conclude at once that the charge carried by 
the one gram of hydrogen ion is not only opposite in sign but 
equal in amount to the electricity required. In general then 
one formula weight of a univalent ion must carry a total charge 
equal to 96,500 coulombs. One formula weight of an ion of any 
valence will carry one, two, three, four, etc., times this charge, 
according to whether its valence is one, two, three, or four, etc. 
If we assume that one formula weight of any one ion represents 
the same number of free ions as a formula weight of any other 
ion (and this is in strict accord with our accepted definition of 
the term formula weight) , we come at once to the conclusion that 
all univalent ions carry equal charges. We call this a unit charge ; 
each bivalent ion carries two unit charges, each trivalent ion carries 
three unit charges, etc. In writing the symbols or formulae of 
free ions it is customary to add one or more + or — signs to 
indicate the number of positive or negative unit electric charges 
carried by the ion, for example, H + , Cu ++ , Al +++ , Cl~, 

so 4 — , po 4 — . 

405. Equilibrium between Molecules and Ions. — The facts 
already studied (396), together with a great volume of other 
evidence which we shall take up in turn, led the Swedish chemist 
Svante Arrhenius to the conclusion that in concentrated solu- 
tions of acids, bases, and salts a considerable part of the dissolved 
substance is present as molecules; but that as the solution is 
diluted, more and more of the molecules dissociate into free ions, 
until in very dilute solutions (at least in many cases) the disso- 
ciation is nearly complete. On the other hand, when a dilute 
solution is evaporated the ions undoubtedly gradually unite 
to form molecules, until, when complete dryness is reached, only 
molecules are present. In any given solution a state of equi- 
librium exists between molecules and ions, as represented in the 
case of common salt by the equation 

NaCl^Na++Cl-. 



The Ionic Hypothesis 



243 



At a definite concentration a definite proportion of the salt will 
be present as ions; this proportion we call the fraction ionized 
or the degree of ionization. We shall next take up the important 
problem of determining the fraction ionized for any solution of 
an electrolyte. Since we believe that the current in a solution 
is carried by the ions present, the ability of a solution to conduct 
a current, or, briefly, its conductivity, must be an indication of 
the extent to which its molecules are dissociated into ions. 

406. Effect of Dilution on Conductivity. — We have already 
learned from the color changes produced by diluting solutions 
that ionization is promoted by dilution. Let us now consider 
the question, What influence, if any, does the volume of water 
in which a given quantity of an acid, base, or salt is dissolved 



H:0 



HCI 



li 



£& 



g&J 



Fig. 51 

have on its electrical conductivity? We may study this ques- 
tion experimentally by means of the apparatus shown in Fig. 5 1 . 
The rectangular glass vessel of about 1 liter capacity is provided 
with two large copper electrodes, as shown in the figure. The 
vessel is first filled about three-fourths full of distilled water, 
and the electrical connections are made. No appreciable current 
passes. Next about 200 c.c. of concentrated hydrochloric acid 
are introduced below the water, without mixing, in such a way 
as to form a separate layer. This may be done by the use of a 
dropping funnel, the stem of which reaches the bottom of the 
vessel. 

The vessel now contains two distinct layers — a lower layer 
of concentrated hydrochloric acid and an upper layer of water. 
The galvanometer indicates that a considerable current is 
passing, and we conclude that this is all passing through the 
acid in the lower layer and not through the upper water layer. 



244 Introduction to General Chemistry 

If next we mix the acid and water thoroughly and so dissolve 
the acid in a much larger volume of water, we note that a large 
increase in the current takes place. This leads us to conclude 
that the conductivity of the hydrochloric acid present is greater 
when it is dissolved in the larger volume of water. We may now 
ask, however, whether there is a limit to the increase in conduc- 
tivity when a given amount of acid is dissolved in larger and 
larger volumes of water, the conductivity being measured under 
such conditions that the solution is all contained between parallel 
plates at a fixed distance apart. 

If, in the experiment described, the vessel were much deeper, 
but otherwise the same, and the electrodes extended all the way 
up the sides as before, it would be found that a given amount 
of hydrochloric acid diluted with double the amount of water 
used in the first experiment would show appreciably greater 
conductivity than in the first case. Or if the acid were diluted 
with three, or four, or still more times as much water, greater 
and greater conductivity would have been observed; but with 
increasing dilution the increase in conductivity would become 
smaller each time more water was added, so that finally a maxi- 
mum conductivity would be reached. Beyond this limit further 
dilution would cause no increase in conductivity. 

These same experiments could be repeated with many other 
electrolytes with similar results. 

407. Definition of Molecular Conductivity. — If one formula 
weight (called also one gram molecular weight) of an acid, base, 
or salt is contained in a solution which is wholly included between 
two parallel electrodes 1 cm. apart, we call the electrical con- 
ductivity of this solution its molecular conductivity. To find 
the molecular conductivity experimentally we measure its 
electrical resistance in ohms. The reciprocal of the resistance 
so found is by definition the molecular conductivity. The con- 
clusions of the paragraph on the effect of dilution on conductivity 
may now be summarized as follows: The molecular conductivity 
of every electrolyte increases as its solution is diluted and finally 
attains a maximum which has a definite numerical value for each 
substance (the temperature being fixed). Table XVI shows the 



The Ionic Hypothesis 245 

change of molecular conductivity of hydrochloric acid as the 

volume in which one formula weight of acid is contained is 

increased. 

TABLE XVI 

The Molecular Conductivity of Hydrochloric 
actd at 1 8° (noyes and cooper) 

Volume of Solution Molecular 

in Liters Conductivity 

IO 351 

12.5 353 

100 368 

5oo 373 

2,000 375 

Maximum 379 

408. Determination of the Degree of Ionization — When a 
solution of a substance is so dilute that it has its maximum molec- 
ular conductivity, it is assumed that all of its molecules have 
dissociated into ions. In a more concentrated solution, for which 
the molecular conductivity is less than the maximum, the frac- 
tion which its observed molecular conductivity forms of its 
maximum molecular conductivity is consequently equal to the 
fraction which the number of ions present in that particular 
solution form of the total number of ions in the completely 
dissociated (completely ionized) solution of the same quantity 
of that substance. This fraction is therefore the fraction ionized 
or the degree of ionization. Thus for decinormal hydrochloric 
acid the degree of ionization is 351-^-379 — 93 per cent. 

This method of determining the degree of ionization was 
proposed by Arrhenius in the year 1887. His reasoning ran 
thus: The passage of a current through a solution is accom- 
plished by the migration of positive ions in one direction and 
negative ions in the other. These transport electricity through 
the solution between the electrodes. Since the molecular con- 
ductivity of a substance is the measure of its rate of transporting 
electricity, it is plain that the molecular conductivity will depend 
on the number of ions present, the charge on each, and the velocity 
of migration. Now under the conditions used in measuring 
resistance, and therefore also of measuring molecular con- 
ductivity, the velocity of migration of its ions will be the same 



246 Introduction to General Chemistry 

for all concentrations of solutions of a given substance (except 
for very concentrated solutions). The charges of the individual 
ions of a given substance are also the same, whether the solution 
is dilute or concentrated. Therefore the molecular conductivities 
of solutions of a given substance are directly proportional to the 
numbers of ions present. Consequently the ratio of the molecu- 
lar conductivity for a given concentration to the maximum mo- 
lecular conductivity for this substance is the fraction ionized, since 
it is assumed that a very dilute solution having maximum molec- 
ular conductivity is completely ionized. 

409. Results of Determination. — The degree of ionization of 
some common electrolytes is shown in Table XVII. 

410. Discussion of Table XVII. — A study of the table leads 
to the very important generalization that in solutions of most 
salts a large percentage of the substance is in the form of ions; in 
consequence, we say that such solutions are highly ionized. It 
also appears that dilute solutions of hydrochloric and nitric acid 
are even more highly ionized than salt solutions of like concentra- 
tion. On the other hand, decinormal acetic acid is only 1 . 3 per 
cent ionized, while the degree of ionization of decinormal carbonic 
acid is very much less, namely 0.17 per cent. In general, the 
extent to which acids are ionized, in solutions of equal concen- 
tration, varies enormously. Bases also differ greatly in their de- 
grees of ionization. For example, decinormal sodium hydroxide 
is ionized 90 per cent, while the same concentration of ammonium 
hydroxide is only 1 . 3 per cent ionized. We have already learned 
that every substance is more highly ionized in dilute than in 
more concentrated solutions. The percentage of ionization of 
a substance as shown in the table applies only to the indicated 
concentration and temperature. 

411. Resume of the Ionic Hypothesis. — We have already 
developed enough of the ionic hypothesis to go far into the under- 
standing of double decomposition reactions. Let us therefore 
review in brief the ideas already brought out, and then, after 
a short critical survey of the fundamental assumptions, proceed 
in chapters xviii and xix to the application of the hypothesis to 
practical examples. 






The Ionic Hypothesis 



247 



According to the ionic hypothesis, as soon as an acid base or 
salt is dissolved in water it is immediately dissociated to some 
extent into ions which prove to be the parts of those substances 
which we have found active in double decomposition. The basic 

TABLE XVII 

Values of the Degree of Ionization of Some Common Electrolytes in 

Water Solution at i8° 
(Degree of ionization at the normality indicated at the head of the column) 



0.05 



Salts: 

NaCl 

KC1 

KBr 

Kl 

NaN0 3 

KNO3.. 

AgN0 3 

KCIO3. 

BaCl 2 

CaCL 

MgCl 2 

PbCl 2 

Sr(N0 3 ) 2 

Ba(N0 3 ) 2 

K 2 S0 4 

Ag 2 S0 4 

MgS0 4 

ZnS0 4 

CuS0 4 

Bases: 

NaOH 

Ba(OH) 2 

NH 4 OH 

Acids: 

HC1 

HNO3 

HC 2 xi30 2 

H 3 P0 4 = H++H 2 P0 4 

H 2 S0 4 = 2 H++S0 4 - 

H 2 C0 3 = H++HC0 3 



81 

87 
86 

87 
84 
67 
63 
63 

96 

93 
04 

97 

97 

042 

60 

64 

005 



93 
86 
017 



94 

94 

020 

36 

38 

002 



45 
4i 
40 

90 
81 
013 

93 

93 

013 

29 

3i 

0017 



0.74 

74 



5i 



59 



£art carries a positive charge and the acidic part a negative 
charge. The charge carried by any univalent ion is called a 
unit charge; ions having greater valence carry as many unit 
charges as they have valence. Since the solution of any elec- 
trolyte is always electrically neutral, the total quantity of posi- 
tive electricity carried by the positive ions equals the total 



248 Introduction to General Chemistry 

quantity of negative electricity carried by the negative ions. 
The more dilute the solution the greater is the proportion of the 
electrolyte transformed into ions, or, in other words, the greater 
is its degree of ionization. 

When we put two electrodes, connected with dry cells or 
other source of current, into a solution of an electrolyte, the 
current is found to flow in the outside circuit, because of the 
discharging of the charge on the electrodes, due to the arrival 
at their surface of oppositely charged ions from the solution. 
The ions in the solution move up to the plates because each ion 
is attracted toward one plate and repelled from the other, owing 
to the fact that it also carries a charge of electricity. The sign 
of the charge on the ion determines the direction of the latter's 
movement. On coming in contact with the electrodes the ions 
become discharged by having their charges electrically neutral- 
ized by equal amounts of the opposite kind of electricity furnished 
by the electrode. Metallic ions after discharge are either de- 
posited as metallic platings on the cathode or react with water 
to form hydroxides and hydrogen. Non-metallic ions such as 
hydrogen, oxygen, and chlorine are released as single atoms which 
then unite to form diatomic molecules of the gases H 2 , 2 , Cl 2 , 
etc. Nitrate and sulfate ions never remain discharged at the 
electrodes, but instead we find there the products of their 
reactions with water — nitric acid and oxygen, and sulfuric acid 
and oxygen, respectively. 

If a given quantity of electrolyte is kept between plates 
which are parallel to each other and carry a constant charge per 
unit area, at different dilutions the conductivity of this elec- 
trolyte will vary in proportion to the number of free ions present. 
As a consequence the proportion of an electrolyte which has 
been transformed into ions at any dilution can be determined 
by dividing the conductivity at the dilution in question by the* 
maximum conductivity found after continuous dilution of this 
same quantity of electrolyte, provided the two measurements 
are made in the manner described (408) . Values so determined 
show us that as a rule all salts are highly ionized substances, 
but acids and bases have very different degrees of ionization, 



The Ionic Hypothesis 249 

some being even more highly ionized than salts, but others 
being very little ionized indeed. 

412. Criticisms of the Ionic Hypothesis. — The idea that ions 
exist in solution as independent chemical substances has come 
to be known as the Ionic Hypothesis. It will be surprising if 
the student who learns of this hypothesis for the first time and 
thinks critically about the matter is not ready to offer at once 
several good reasons for doubting the truth of the conclusions. 
In the first place, the hypothesis seems to assume that in a solu- 
tion of common salt, for example, a large part of its elements 
are present in a free state. Now the student who knows any- 
thing of the properties of metallic sodium and of chlorine gas 
will find it hard to believe that either of these elements can be 
present in a salt solution; because sodium reacts violently with 
water, forming sodium hydroxide and hydrogen, and chlorine 
has a horrible smell and a yellow color. Plainly there is some- 
thing incompatible with the obvious facts in the statement that a 
solution of salt contains free sodium and free chlorine. 

A closer study of the hypothesis shows, however, that it is 
not assumed that the elements sodium and chlorine are present 
as ordinary atoms, for each atom is said to be electrically charged. 
Those who uphold the hypothesis will point out that a charged 
brass ball has very different properties from the same ball if 
uncharged. True, say the critics, but even a charged brass ball 
is still a brass ball; to which the opponents reply that the 
quantity of electricity on the ball is a matter of enormous 
importance. 

If then the ions are so highly charged, why do the positive 
ions not unite with and so electrically discharge the negative 
ions, since the solution is a conductor? It may be said in reply 
that it is assumed that ions of unlike sign are constantly uniting, 
at a rate just equal to the rate of dissociation, with the result 
that a state of equilibrium is produced. 

In spite of the foregoing criticisms and many others the ionic 
hypothesis with all its apparent inconsistencies has proved itself 
highly useful in explaining and correlating many facts and 
phenomena. 



250 Introduction to General Chemistry 

Before passing final judgment on this remarkable hypothesis 
it will be better to consider its further applications and then, 
in chapter xx, to take up the matter again in the light of newer 
discoveries, which have led to essential modifications of the 
views as originally proposed by Arrhenius. 

Finally it may be urged that Arrhenius himself was not 
certain of the. truth of his theory until he became acquainted 
with the wonderful work of Van't Hoff on the so-called osmotic 
pressures of dissolved substance (chap, xxvii). This work 
will be discussed as soon as we have progressed far enough to 
understand and interpret the experiments which we must then 
study. 



CHAPTER XVIII 
APPLICATIONS OF THE IONIC HYPOTHESIS 

413. Double Decomposition. — In the foregoing chapter it 
was pointed out that the probable cause of the union of two 
unlike ions is the attraction of their unlike electric charges. In 
general, every kind of positive ion can unite with any kind of 
negative ion. Therefore, if any two electrolytes (provided they 
have no ion in common) are mixed in solution, at least some 
double decomposition must take place, simply because new 
combinations of positive and negative ions are made possible. 
Let us first consider the important case in which the two starting 
materials, as well as the two products of the reaction, are easily 
soluble and highly ionized. 

414. Class I. Equilibrium between Four Easily Soluble and 
Highly Ionized Electrolytes. — If dilute solutions of two imaginary 
electrolytes A B and CD, which ionize thus 

AB^A++B~, I 

CD±^C++D~, II 

are mixed, we may predict, without knowing anything more 
about these substances, that the following reactions are possible, 

A++D~±^AD, III 

C++B-±;CB, IV 

and that a double decomposition reaction, 
AB+CD±+AD+CB, 

will take place to a greater or less extent. Since all four of the 
substances AB, CD, AD, and CB are assumed to be highly 
ionized, it is plain that the mixed solution will contain chiefly the 
four kinds of ions, A + , B~, C + , and D~ , and relatively few 
molecules. Since each of the four kinds of molecules present 
must be in equilibrium with its own two kinds of ions, the four 



252 Introduction to General Chemistry 

equilibrium reactions (marked I, II, III, and IV) must be 
interrelated in the manner shown by the following arrangement 
of equations I, II, III, and IV: 

III IV 

AB^A++B~ I 

CD^D~ + C+ II 

It It 
AD CB 

Equations I and II read horizontally, while III and IV read 

vertically. We may call this the compound equation of the 

reaction 

AB+CD^AD+CB. 

The compound equation shows that the four molecular sub- 
stances are in equilibrium with each other because each molec- 
ular substance is in direct equilibrium with its own pair of ions- 
Now if all of the four molecular substances are assumed to have 
exactly equal tendencies to ionize, then we must conclude that 
for the condition of equilibrium equal numbers of the four kinds 
of molecules will be present, if we have taken equivalent amounts 
of substances. We may summarize Class I as follows: If both 
starting substances and both products of a double decomposition 
reaction AB-\-CD*=> AD-{-CB are easily soluble and highly 
and equally ionizable, an equilibrium mixture will result in 
which (1) most of the dissolved material is present as free ions, 
(2) little of the material is present as molecules, and (3) if 
equivalent amounts are taken the four kinds of molecules are 
present in equal numbers. 

415. An Example of Class I. The Reaction between Ferric 
Chloride and Ammonium Sulfocyanate.— The reaction 
FeCl 3 +3NH 4 NCS±>Fe(NCS) 3 +3NH 4 Cl, 

studied earlier (280), is a good illustration of Class I, since all 
four salts are easily soluble and highly ionized. It was shown 
by experiment that this reaction does not take place completely, 
but that it reaches equilibrium while there is still much of the 
material not converted into ferric sulfocyanate and ammonium 
chloride. 




Applications of the Ionic Hypothesis 253 

416. Other Examples of Class I. — Numerous additional 
examples of Class I might be given. The following will serve 
as illustrations: 

KCl+NaN0 3 ±5KN0 3 +NaCl, 

Na 2 S0 4 + 2KNO3 ±5 K 2 S0 4 + 2 NaN0 3 , 

K 2 C0 3 +Na 2 S0 4 ±5K 2 S0 4 +Na 2 C0 3 , 

NaN0 3 +HCl±5NaCl+HN0 3 . 

In each case the mixed solution contains largely the four kinds 
of ions, together with small proportions of the four kinds of 
molecules in approximately equivalent amounts. In Class I 
the two substances taken react only partially, and there- 
fore the reaction is always incomplete. 

417. A Graphic Method of Representing Degrees of 
Ionization. — An acid, base, or salt, not in solution, exists FlG 
wholly in the form of molecules (no ions are present). 

We may represent such an un-ionized substance by a cross- 
hatched circle, Fig. 52. When this substance, whose formula we 
may call AB, is dissolved in water it will partially ionize, thus: 

AB^A++B~. 

This condition is represented by Fig. 53. Let us suppose that 
the solution is 80 per cent ionized ; then 20 per cent is present as 
un-ionized molecules. In Fig. 53 the left-hand circle has a 

cross-hatched sector winch is 
ab ^MiHrrKA + /*^^\ b- just 20 per cent of the area 

of the whole circle. This will 

represent the fact that 20 
Fig. 53 per cent of the substance is 

present as un-ionized mole- 
cules. The middle circle, of which 80 per cent is shaded with 
vertical lines, will represent the fact that 80 per cent of the 
total A radical is in the form of free positive ions. In similar 
fashion the left-hand circle shows that 80 per cent of the 
total B radical is in the form of free negative ions. Further- 
more, if we take the area of the circle, Fig. 52, as propor- 
tional to the whole number of molecules in one formula weight 
of the substance before it is dissolved, then the area of the 




254 



Introduction to General Chemistry 




CuSO. 




cross-hatched sector of the left-hand circle of Fig. 53 will be pro- 
portional to the number of un-ionized molecules in one formula 
weight of the dissolved substance. Since each AB molecule, 

when it ionizes, gives one A + 
ion and one B~ ion, the areas 
of the shaded portions of the 
middle and right-hand circles 
will be directly proportional to 
the numbers of A + and B~ 
ions respectively. By means 
of a figure like Fig. 53 the 
relative concentrations of ions 
and molecules of a dissolved 
electrolyte can be seen at a 
glance. By way of further 
illustration the condition of 



+ / 



Cu ++ 



^ N CuSO. 




so.— 



Fig. 54 



normal, one- tenth-normal, and one one-hundredth-normal copper 
sulfate solution is shown in Fig. 54. 

418. Graphic Representation of Class I. — Let us now turn 
to the graphic representation of a double decomposition reaction 
of the type just studied under Class I, where all four substances 
concerned are easily soluble and highly ionized. We again repre- 
sent the reaction by 

AB+CD^AD+CB. 

Figure 55 shows the condition of solutions of AB and CD 
before they are mixed, on the supposition that each is 85 per cent 
ionized in N/ 10 solution. When 
equal volumes of the two N/10 
solutions are mixed, the reaction 
represented by the compound 
equation 

AB^A++B~ 

CD^D~+C+ 



CD 




AD CB 



Fig. 55 



takes place and very quickly reaches the condition of equilib- 
rium shown graphically in Fig. 56, in which the proportions of 




Applications of the Ionic Hypothesis 255 

molecules and ions have been calculated on the additional 

assumption that AD and CB both have the same tendency to 

ionize as have AB and CD (when each is separately dissolved 

in water). Comparison of Figs. 55 and 56 shows us that the 

areas representing the numbers 

of molecules and ions of the S ~^ AB 

materials taken are not greatly ^ > > 

changed as the result of the 

mixing. Consequently we say /"\cd 

that the reaction is incomplete. 

All examples of Class I would be 

represented by similar graphs. 

419. A Second Type of §&> » jL, \ 
Double Decomposition: Class V _ ^ ' 

II. — Class II will comprise Fig. 56 

double decomposition reactions 

in which two easily soluble and highly ionized substances give 

two easily soluble products, one of which is highly ionized, the 

other little ionized. The simplest example of Class II is found 

in a neutralization reaction such as 

HCl+NaOH^NaCl+H 2 0, 

since all the substances except the water are highly ionized. 

420. The Ionization of Water. — The ionization of water may 
be determined from conductivity measurements, for though it is 
a very much poorer conductor of the current than is a salt solu- 
tion, still, as we have already said, it conducts much better 
than glass or hard rubber. According to the ionic theory it is 

ionized thus: 

H 2 O^H++OH". 

In one liter of pure water there is present about one ten-millionth 
of a gram of ionic hydrogen and the equivalent amount of 
hydroxyl. 

If then we attempt to represent the proportion of ions in pure 
water by a graphic scheme, a single dot in the center of an other- 
wise empty circle would have too large an area to represent 
correctly the proportion of ions present if the rest of the circle 




256 Introduction to General Chemistry 

represented the molecules of water. In cases of this kind we 
shall use a single dotted radius to indicate that the number of 
ions is too small to be accurately represented. The graph of 

water will then be that shown in 

- i — \ { — : Fig. 57. 

V_.-' v_,*' That there are so few ions 

FlG present in a liter of pure water 

means that the ionic equilibrium 
is established only when all but a minute fraction of the total 
material is in the form of water molecules. Accordingly, when 
hydroxyl and hydrogen ions are brought together in solution 
we must expect them to combine almost completely to form 
molecules. 

421. Neutralization. — If we mix equivalent amounts of 
solutions of HO and NaOH the resulting reaction may be 
represented as follows: 

HC1^H++C1~ 
NaOH^OH-+Na+ 

It It 

H 2 NaCL 

The H + and OH - ions present unite almost completely to form 

molecules of H 2 0. The removal of H + ions causes a shift of the 

reaction 

HC1±5H++C1- 

to.the right, and as the H + ions produced in this way are almost 
immediately taken up by new OH~ ions formed by a shift to the 
right of the reaction 

NaOH^OH"+Na+, 

the final result is the practically complete dissociation of both 
HC1 and NaOH molecules and therefore the disappearance of 
these substances. Molecules of H 2 0, once formed, dissociate 
very little into H + and OH" ions, and so the final equilibrium 
solution will contain no more free H + and OH~ ions than an equal 
volume of pure water. The Na + and Cl~ ions unite partially to 

form molecules 

Na++Cl-^NaCl, 



Applications of the Ionic Hypothesis 



257 



but this reaction does not proceed far in dilute solution, as com- 
mon salt is a highly ionized substance. In fact, the solution 
resulting from the neutralization of HC1 by.NaOH is exactly 
the same as, and differs in no way from, that made by dissolving 
common salt in water to produce a solution of equal concentra- 
tion. All the facts just stated are shown by a comparison of the 
two graphs, Figs. 58 and 59. Thus it can be seen (Fig. 59) that 
the areas representing the numbers of molecules of HO and 
NaOH respectively have been reduced to negligible dimensions; 
the same is also true of the areas for H + and OH~ ions. But 



sHCI 

\ 

! 



NaOH 





^"~~"XNaOH 




Fig. 58 



Fig. 59 



the Cl~ and Na + ion areas are not greatly changed in the second 
graph. Compared with these areas, that representing NaCl mole- 
cules is small. The circle representing the number of molecules 
of H 2 is completely shaded, thus showing that the yield of 
molecular H 2 is practically 100 per cent. 

422. The Simplified Equation of Neutralization. — To sum 
up the matter, it may be said that acids and bases neutralize 
one another because of the tendency of H + and OH~ ions to unite 
almost completely to form water. This almost complete union 
of H + and OH~ ions takes place because H 2 is but very slightly 
ionized. In a very dilute solution, where the acid and base 
taken are almost completely ionized at the instant of mixing, 
the principal change that takes place is the union of H + and 
OH~ ions to form H 2 molecules, since in the very dilute solu- 
tion the Na + and Cl~ ions remain largely uncombined. We may 



258 Introduction to General Chemistry 

therefore write as the simplified equation of neutralization in 
dilute solution 

H++OH-^H 2 0. 

423. Experimental Confirmation of the Theory of Neutraliza- 
tion. — The process of neutralization can be followed experi- 
mentally with the help of an apparatus somewhat like that shown 
in Fig. 51 (406); but having a small electric lamp in the place 
of the galvanometer. The solution layers in the cell shown 
in the figure are made by first putting into the cell a layer of 
one-tenth-normal hydrochloric acid, and then allowing an equal 
layer of sodium hydroxide to run under this first layer by intro- 
ducing it at the bottom of the cell through a dropping funnel. 
As represented in the figure the two parallel electrodes are in 
contact with the two layers, which can be seen very nicely if a 
little litmus is put into the acid and base respectively before the 
layers are made. If now the key is closed the current flows 
through both layers, and the lamp glows. Hydrogen and sodium 
ions are arriving at one electrode, and chlorine and hydroxyl 
ions are arriving at the other. Of these ions the hydrogen and 
hydroxyl travel much more rapidly under the attraction from 
a given charge per unit area of the electrode, and so they are 
neutralizing their charges on the plates more quickly than are 
the other ions. As a result most of the current passing in the 
outside circuit is due to their discharge. If the two layers of 
acid and base are next mixed, the lamp no longer glows. Half 
of the carriers of the current and the most efficient ones have 
been used to form water molecules, and in the cell there remains 
only the slow-moving sodium ions and chlorine ions. If the 
acid and base in the respective layers were not quite equivalent 
in amount, a slight excess of one or the other will be shown by 
the litmus color, but the important part of the experiment, the 
serious loss of ions, will still be unmistakable from the great 
decrease in the conductivity of the solution between the plates. 

424. A Second Example of Class II: Action of HC1 on 
Sodium Acetate. — It will be recalled that acetic acid, HC 2 H 3 2 , 
neutralizes NaOH, forming sodium acetate, thus: 
HC 2 H 3 2 +NaOH^NaC 2 H 3 2 +H 2 0. 



Applications of the Ionic Hypothesis 259 

Acetic acid is a monobasic acid, only one of the four hydrogen 
atoms of each molecule being ionizable: 

HC 2 H 3 2 ^H++C 2 H 3 2 -. 

This acid is but little ionized in normal solution, the degree of 
ionization being only 0.4 per cent. On the other hand, solu- 
tions of its salts, like NaC 2 H 3 2 , are highly ionized: 

NaC 2 H 3 2 ±5Na++C 2 H 3 2 -. 

If we mix equivalent amounts of HC1 and NaC 2 H 3 2 in 
solution we cannot see that any chemical change occurs; but 
that a reaction has occurred we may show convincingly with the 
help of the electrolytic cell, which is used to discover the change 
in conductivity during neutralization. In the lower layer this 
time we shall have sodium acetate and in the upper hydrochloric 
acid. As before, the lamp glows — both solutions are good con- 
ductors; the first by means of Na + and C 2 H 3 2 ~~ ions, the 
second by means of H + and Cl~ ions. When we mix the two 
layers the decrease in brightness of the lamp shows that the 
conductivity has dropped off greatly, thus proving that many of 
the ions have been changed into non-conducting molecules. 
The compound equation is 

HC1^H++C1~ 

NaC 2 H 3 2 ^C 2 H 3 2 ~+Na+ 

It It 

HC 2 H 3 2 NaCl 

The graphs are shown in Figs. 60 and 61. Since of the four sub- 
stances concerned all but the acetic acid are highly ionized, while 
the latter is but little ionized, the reaction falls under Class II. 
When HC1 and NaC 2 H 3 2 solutions are mixed, the H + and 
C 2 H 3 2 ~ ions will unite far more completely than will any other 
pair of ions, and at the same time the molecules of HC1 and 
NaC 2 H 3 2 will continue to ionize until but very few remain 
(Fig. 61). Also Na + and Cl~ ions will unite partially to form 
molecules of NaCl. Therefore the equilibrium mixture will 
contain largely free acetic acid, for the most part un-ionized, 



260 



Introduction to General Chemistry 



together with common salt and its ions. Very little HC1 and 
NaC 2 H 3 2 will be present. 

425. Comparison of the First and Second Examples of Class 
II. — Fig. 60 shows the conditions of the solutions of hydrochloric 
acid and of sodium acetate before they are mixed, while Fig. 61 
shows the condition of the equilibrium mixture. These figures 
are almost a reproduction of Figs. 58 and 59, representing neutral- 
ization. In place of NaOH we have in the second case NaC 2 H 3 2 , 
which is also highly ionized; and in place of water we have 
HC 2 H 3 2 , which, like water, is but little ionized. In Fig. 61 the 
circle representing molecular acetic acid is nearly completely 

Ghci ''~r.ii* 

» • ; ,.7 < 



S HCI 
1 



^^_,' 



, NaCH.O* 



^_^ 





NaC 2 H 3 2 ,'~^ N C 2 H 3 s 




n n 

k HCH,0, 




Nad 



Fig. 60 



Fig. 6] 



cross-hatched, showing that the yield of this substance is nearly 
100 per cent. Although the reactions represented by Figs. 59 
and 61 are so nearly alike, there is a small difference due to the 
fact that acetic acid is ionized more than water. In consequence 
the formation of molecular acetic acid falls short of 100 per cent 
by a small fraction of 1 per cent. 

426. A Third Example of Class II: Action of NaOH on 
NH 4 C1. — Another important example of Class II is found in 
the action of sodium hydroxide and ammonium chloride. The 
addition of dilute NaOH to a solution of NH 4 C1 does not produce 
any visible effect; but evidence that the reaction 

NaOH+NH 4 Cl^NaCl+NH 4 OH 

takes place may be obtained in two ways: first, by finding a 
great decrease in conductivity on mixing superimposed layers 






Applications of the Ionic Hypothesis 



261 



of the two solutions; and secondly, by noting the odor of 
ammonia given off by reason of the partial dissociation of the 
NH 4 OH present in the solution 

NH 4 OH±5NH 3 +H 2 0. 

The compound equation of the reaction follows : 

NaOH±5Na++OH- 

NH 4 C1±5C1-+NH 4 + 

If It 

NaCl NH 4 OH. 

Since sodium hydroxide and ammonium chloride are highly 
ionized, and ammonium hydroxide is little ionized, this reaction 
is completely analogous to that between HC1 and NaC 2 H 3 2 : 

HCl+NaC 2 H 3 O^NaCl+HC 2 H 3 2 . 

Each reaction takes place nearly completely from left to right 
because one product is but little ionized. The graphs for 
this reaction, Figs. 62 and 63, are closely similar to those for 



'""-xNaOH 



Na + • " ~ - . 0H- 





Fig. 6: 



Fig. 63 



neutralization, Figs. 58 and 59, and for the reaction between 
HC1 and NaC 2 H 3 2 , Figs. 60 and 61. 

427. Summary of Class II Reactions. — As we have pointed 
out, all these reactions are alike, in that two highly ionized 
electrolytes react to form one highly ionized electrolyte and one 
little ionized electrolyte. Invariably reactions of this class are 
nearly complete. The smaller the degree of ionization of the 
little ionized product, the more completely the reaction takes. 



262 Introduction to General Chemistry 

In the resulting mixture the little ionized substance is present, 
of course almost wholly in the molecular form. 

428. Strength of Acids. — Since all salts are highly ionized, 
the reaction between any highly ionized acid and a salt of a 
little ionized acid must belong to Class II. We may therefore 
predict that, as in the second example studied, such reactions 
will give nearly 100 per cent yields of their products, and that 
in the resulting solution there will be present the little ionized 
acid instead of the highly ionized acid originally used. The 
highly ionized acid may be said to have displaced the little 
ionized acid from its salt. As a result, we may call the former 
a strong acid and the latter a weak acid, and may say that a 
strong acid always displaces a weak acid from its salts. 

429. Strength of Bases. — Just as we call a highly ionized 
acid a strong acid and a little ionized acid a weak acid, so we 
may call a highly ionized base a strong base and a little ionized 
base a weak base. Since all reactions between strong bases and 
the salts of weak bases (see third example, 426) are examples of 
Class II, we can predict that the yield of weak base and salt of 
the strong acid will be nearly 100 per cent. In other words, a 
strong base will always displace a weak base from its salt. 

430. Two Useful Laws. — -The foregoing law and that given 
in the paragraph on the strength of acids (428) have been of 
very great practical convenience to chemists. These laws fail 
only when the salt of the weak acid is little ionized, a case so 
rare that the usefulness of the rules is virtually unaffected. The 
laws are of course only special cases of the fundamental one that 
if two highly ionized substances react to form one little ionized sub- 
stance and one highly ionized substance, the reaction will be nearly 
complete. 

431. Suppression of the Ionization of a Weak Acid or a 
Weak Base. — Since the strength of an acid or a base is deter- 
mined by its tendency to ionize, any factor that has an influence 
on this tendency will affect the strength or weakness of the acid 
or base. We must now consider this important subject and will 
begin by studying the action of NH 4 C1 on a solution of the weak 
base NH 4 OH. 



Applications of the Ionic Hypothesis 



263 



If we add a little phenolphthalein to very dilute NH 4 OH a 
bright, red-colored solution results. This shows that the solu- 
tion is alkaline, and therefore that it contains an abundance of 
OH~ ions. Upon addition of a little NH 4 C1 to this red solution 
the color disappears almost completely. This proves that the 
number of OH~ ions present has been very greatly decreased. In 
order to understand how this has happened, we must consider 
the matter from the standpoint of ionic equilibrium. A solu- 
tion of NH 4 OH is ionized to a small extent, thus: 

NH 4 OH^NH 4 ++OH". 

Ammonium chloride, on the other hand, is very highly ionized: 

NH 4 C1^NH 4 ++C1-. 

If then we add an equivalent amount of NH 4 C1 to a dilute solu- 
tion of NH 4 OH, the number of NH 4 + ions per cubic centimeter 
will be increased many fold. The OH~ ions present will therefore 
collide with NH 4 + ions and combine with them far more fre- 
quently than before. Since the rate of dissociation of NH 4 OH 
molecules into ions is not affected by the presence of the NH 4 C1, 
this increased rate of union of NH 4 + and OH~ ions causes a 
great shift to the left of the equilibrium 

NH 4 OH±5NH 4 ++OH-. 

For example, it has been found, by methods that we need not 
consider here, that the addition of 1 g. of NH 4 C1 to 100 c.c. of 




NH.OH 



• - ^ NH.CI 



- - N NH, + 

\ 




Fig. 64 




mm 



" - - NHXI 




Fig. 65 



decinormal NH 4 OH will decrease the number of OH~ ions 
present about one hundred fold. In other words, the ionization 
of the base will be decreased one hundred fold (see Figs. 64 and 
65). We may now state the general law of which the case just 



264 Introduction to General Chemistry 

studied is a typical example: The ionization of a weak base is 
greatly suppressed by the addition of a salt of the base. This 
means that a weak base is made still weaker by the addition of its 
soluble salts. 

In a similar manner the ionization of a weak acid is greatly 
suppressed by the addition of any of its soluble salts; that is, a weak 
acid is made still weaker by adding one of its salts. For example, 
the addition of NaC 2 H 3 2 to a red solution of acetic acid, 
HC 2 H 3 2 , containing litmus changes the color from red to 
purple, thus showing a great decrease in the number of H + ions, 
and therefore a great decrease in ionization of the acid. 

432. The Common Ion Law. — A base and any one of its 
salts must of necessity have one ion in common. (The NH 4 + 
ion is common to XH 4 OH and XH 4 C1.) An acid also must have 
one ion in common with any of its salts. We may therefore 
state the principle of the foregoing laws as follows : Suppression 
of the ionization of a little ionized substance occurs when we add to 
its solution a highly ionized substance having a common ion. This 
is the Common Ion Law, a very important generalization. The 
examples already cited are by no means the only ones of impor- 
tance. For example, it is plain that the ionization of NH 4 OH 
must be suppressed by the addition of XaOH or KOH because 
of the increase in concentration of the common OH~ ion; and 
that the ionization of HC 2 H 3 2 must likewise be suppressed by 
the addition of any strong acid like HC1 or HX0 3 . The effect 
of a highly ionized substance on the ionization of another highly 
ionized substance having one ion in common is of the same type 
but very much smaller in degree than when the second substance 
is slightly ionized. 

We shall next consider the application of the Common Ion 
Law to solutions of acids and bases and thus obtain a definition 
of the term neutrality. 

433. Neutrality Denned. — We have already learned (420) 
that water is slightly ionized, thus, 

H 2 O^H++OH~. 

Each cubic centimeter of pure water must therefore contain 
exactly as many H + as OH - ions. Since all acids give H + 



Applications of the Ionic Hypothesis 265 

ions, the addition of an acid to water, in accord with the common 
ion law, will greatly suppress the ionization of water. Therefore 
acid solutions will contain far less OH~ ions per cubic centimeter 
than pure water. In an acid solution the number of H + ions 
greatly exceeds the number of OH~~ ions. The ionization of 
water is also greatly suppressed by the addition of a base, since 
all bases have OH~ ions in common with water. In basic solu- 
tions the number of H + ions per cubic centimeter is far less than 
in pure water and therefore the number of OH - ions greatly 
exceeds the number of H + ions. Since we may consider water 
a typically neutral substance we may define a neutral solution as 
one in which the number of H + ions equals the number of H~ 
ions. Since, as we have already learned, a strong acid completely 
neutralizes a strong base, as for example in the reaction 

HCl+NaOH^NaCl+H 2 0, 

we conclude that in the resulting solution the number of H + ions 
is just equal to the number of OH - ions: this is the criterion 
of complete neutrality. 

434. First Example of Class III: The Action of a Weak Acid 
on a Strong Base. — Under Class III we shall study reactions in 
which one little ionized and one highly ionized substance give 
products, one of which is little ionized, the other highly ionized. 
As the first example we shall study the reaction between little 
ionized acetic acid (a weak acid) and highly ionized sodium 
hydroxide (a strong base). These react thus: 

HC 2 H 3 2 +NaOH^H 2 0+NaC 2 H 3 2 . 

Of the products, water is very slightly ionized, while sodium 
acetate, NaC 2 H 3 2 , is highly ionized. If we mix equal volumes 
of normal solutions of HC 2 H 3 2 and NaOH, that is, if we add 
to the NaOH solution exactly that quantity of acetic acid that 
would neutralize it if the reaction were complete, we hnd that 
the resulting mixture is not neutral but is still alkaline to litmus. 
The fact that the mixture is alkaline means that the number of 
HO~ ions is greater than the number of H + ions present. The 



266 



Introduction to General Chemistry 



cause of this condition is most easily understood by aid of the 
compound equation 

HC 2 H 3 2 ^H++C 2 H 3 2 - 

NaOH±50H-+Na+ 

It If 

H 2 NaC 2 H 3 2 

and Figs. 66 and 67. At the instant of mixing, the solution 
contains an abundance of OH~ ions (Fig. 66); these reduce 
greatly the number of H + ions present by forming H 2 molecules : 

H++0H-±5H 2 0. 

The removal of H + ions disturbs the equilibrium 
HC 2 H 3 2 ^H++C 2 H 3 2 -, 

which shifts greatly to the right, thus producing both H + and 
C 2 H 3 2 ~ ions. While the former unite with OH" almost (but 




.H* .' ~nCjH,0, 




HCH.0. 



,--"-~ v NaOH 




C,H,0r 



Fig. 66 



Fig. 67 



not quite) completely, the latter remain for the larger part free 
in the solution, and by their great tendency to unite again with 
H + ions to form little ionized HC 2 H 3 2 serve still further to 
diminish the number of free H + ions. In the final equilibrium 
mixture, shown in Fig. 67, the number of OH~ ions is greater 
than the number of H + ions. because of the great tendency of the 
latter to unite readily with either C 2 H 3 2 ~ ions or OH~ ions. 
That the OH~ ions get by far the lion's share of the H + ions is 
owing to the fact that water is much less ionized than acetic acid. 



Applications to the Ionic Hypothesis 267 

Since the mixture contains more OH~ than H + ions (see Fig. 
67) it is not neutral but alkaline. 

435. The Action of Water on Sodium Acetate. — In the fore- 
going paragraph we studied the equilibrium 

HC 2 H 3 2 +NaOH^NaC 2 H 3 2 +H 2 0. 

The composition of the equilibrium solution was shown in Fig. 67. 
It must be plain from the deduction of 281, that exactly the same 
equilibrium solution would be obtained if we should dissolve in 
the same quantity of water pure sodium acetate, NaC 2 H 3 2 , in 
exactly the amount that would be produced by the complete 
union of all the HC 2 H 3 2 and NaOH used in the first case. As a 
matter of fact we find that a solution of pure sodium acetate is not 
neutral but alkaline to litmus. The action of H 2 on NaC 2 H 3 2 
takes place thus: the salt first dissolves and at once ionizes 
highly to form many Na + and C 2 H 3 2 ~ ions. Water, although 
but slightly ionized, contains some H + and OH~ ions. Occasional 
collisions of H + and C 2 H 3 2 ~ ions will occur, and part of these 
collisions will result in unions to form HC 2 H 3 2 molecules; and 
as the latter have but little tendency to ionize, the result is a 
great decrease in the number of H + ions present. This in turn 
disturbs the equilibrium 

H 2 O^H++OH-, 

which in consequence shifts to the right and so brings more OH~ 
ions into the solution. A few but not many of the OH~ ions 
unite with Na + ions to form molecules of NaOH, but most of 
the OH - ions remain free, thus producing in the solution a 
decided excess of OH~ ions over H + ions (see Fig. 67), and so 
making the solution alkaline to litmus. Briefly stated, water 
acts on sodium acetate to a small extent, thus, 

NaC 2 H 3 2 +H 2 O^HC 2 H 3 2 +NaOH, 

and since NaOH is highly ionized, while HC 2 H 3 2 is little ionized, 
the reaction of the solution is alkaline. The composition of a 
water solution of sodium acetate is that shown in Fig. 67. 

436. Hydrolysis of Salts. — The soluble salts of all weak acids 
with the strong bases sodium, potassium, calcium, or barium 



268 Introduction to General Chemistry 

hydroxide give alkaline solutions when dissolved in water. In 
every case the reason is the same as that given for the alkaline 
reaction of sodium acetate solution. The effect of water on the 
salt of a weak acid and a strong base is an example of the type 
of reaction called hydrolysis (or also hydrolytic dissociation). 
Hydrolysis may be defined as a double decomposition reaction 
in which water is one of the reacting substances. The solutions 
of salts of all weak acids and strong bases are alkaline in reaction. 
Other things being equal, the weaker the acid from which the salt 
is derived the greater the extent of the hydrolysis; that is, the 
greater the alkalinity of the solution. 

On the other hand, some salts (other than acid salts like 
NaHS0 4 ) give solutions that have an acid reaction (176). 
Among such are the chlorides, sulfates, and nitrates of copper, 
lead, iron, zinc, aluminum, etc. Experiments show that the 
hydroxides of all these elements are weak bases. It would 
therefore seem probable that the acidity of solutions of the salts 
of these bases with strong acids is due to hydrolysis, and that 
the behavior of such salts with water is the counterpart of the 
behavior of salts of weak acids with strong bases. 

437. A Second Example of Class III: The Action of a Strong 
Acid on a Weak Base. — The action of a strong acid on a weak 
base is plainly the reverse of that just discussed: the action of 
water on the salt of a strong acid and weak base. It follows that 
a weak base does not react completely with the theoretical or 
chemically equivalent amount of a strong acid, and in con- 
sequence the resulting mixture is still acid in its reaction. The 
action of HC1 on the weak base NH 4 OH will serve as a simple 

illustration: 

NH 4 OH^OH"+NH 4 + 
HC1^H++C1" 
It It 
H 2 NH 4 C1. 

Comparison of this reaction with that for HC 2 H 3 2 and XaOH 
where we have a weak acid and strong base will bring out com- 
plete analogy. Experiment shows that a solution of XH 4 C1 in 
water is not neutral but slightly acid in reaction. Briefly 



Applications of the Ionic Hypothesis 



269 



stated, NH 4 OH does not completely neutralize an equivalent 
amount of HO because it is a weak base. Conversely, water 
acts on pure NH 4 C1 to form some free HC1 and NH 4 OH. 

438. Class IV: The Action of a Weak Acid and a Weak 
Base. — Under Class IV we shall include reactions between two 
little ionized substances, which give as products one little ionized 
and one highly ionized substance. The only reactions of im- 
portance in Class IV are those between a weak acid and a weak 
base, the products being water and a salt. Acetic acid and 
ammonium hydroxide are both moderately weak (but not 
extremely weak) . They react thus : 

HC 2 H 3 2 +NH 4 OH^NH 4 C 2 H 3 2 +H 2 0. 

The reaction is not complete, as in the case of the action of a 
strong acid and a strong base, but reaches equilibrium when a 
few tenths of 1 per cent of the free un-ionized acid and free 
un-ionized base are still present in the solution. The conditions 
before and after the reactions are shown in Figs. 68 and 69. 



HCAO, 




- C,H,0r 




Fig. 68 



Fig. 69 



If on the other hand the solution is made by dissolving solid 
NH 4 C 2 H 3 2 in water partial hydrolysis takes place, giving a 
mixture the composition of which is also represented by Fig. 69. 

If both acid and base are extremely weak the extent of the 
hydrolysis will be much greater than in the case of NH 4 C,H 3 2 . 
In fact, in such cases hydrolysis may be so nearly complete that 
we may say that extremely weak acids in water solution do not 
form salts with extremely weak bases. 



270 Introduction to General Chemistry. 

439. Heat of Ionization. — The heat liberated or absorbed by 
the complete dissociation into its ions of one formula of a dis- 
solved electrolyte is called its heat of ionization (cf. 366). In 
some cases heat is absorbed, in other cases it is liberated, when 
the substances are ionized, but in the great majority of cases 
the heat of ionization is very small. For practical purposes we 
may say that the heat of ionization of readily ionizable elec- 
trolytes is almost negligible. Little ionized substances often 
have appreciable heats of ionization. This is notably the case 
with water, for which we have the following, 

H++OH-^H 2 0+ 13,700 cal. 

It was stated earlier (362) that the heat of neutralization of 

a strong acid by a strong base is almost the same in all cases, 

namely 13,700 calories. The reason can now be seen. We 

know that in the neutralization of a strong acid by a strong base 

in dilute solution the principal change is the union of H + and OH~ 

ions to form water. In other words, the simplified equation of 

neutralization is 

H++OH-->H 2 0. 

Since strong acids and bases, as well as most salts, have 
negligible heats of ionization; and since, moreover, very little 
dissociation or union of ions, other than H + and OH~, occurs 
in neutralization (421, Fig. 59), the heat produced in the reaction 
is simply that due to the formation of water from its ions. It 
is for this reason that heats of neutralization are practically 
the same for all strong acids and bases: 13,700 cal. for one 
formula weight (18 g.) of water formed. 

The heat of neutralization of ammonium hydroxide by a 
strong acid is 12,300 calories. The difference, 13,700—12,300 = 
1,400 cal., is the heat of ionization of the weak base. 

In reactions between solutions . of two highly ionized salts 
which form by interaction two other highly ionized and easily 
soluble salts no appreciable heat change is observed. This is 
because in such reactions very little change takes place (418, 
Fig. 55), and such changes as do occur are accompanied by nearly 
negligible heats of ionization. 



Applications of the Ionic Hypothesis 271 

440. Indicators. — In addition to litmus, which is used so 
often to indicate the acidity or alkalinity of solutions, a number of 
other colored substances are also employed. These are called 
indicators. The more important indicators besides litmus are 
phenolphthalein and methyl orange. The former is a colorless 
substance which gives a bright red solution with alkalies. 
Methyl orange is orange color in neutral solution, pink in acid, 
and yellow in alkaline solution. In general, indicators are very 
complex chemical substances whose formulae need not be con- 
sidered at present. 

Since acid solutions always contain H + ions and alkaline 
solutions OH~ ions, we may say that an indicator is a substance 
which has one color in the presence 0} an excess of H + ions and 
a different color in the presence of an excess of OH~ ions. We 
might expect that every indicator would show its transition 
shade of color in an exactly neutral solution; that is, in a solu- 
tion where the number of H + ions equals the number of OH~ 
ions. This, however, is not the case. In other words, most 
indicators do not indicate perfect neutrality. Litmus is a nearly 
perfect indicator, but phenolphthalein shows a change of color 
when the number of OH~ ions equals eighty times the number 
of H + ions; that is, if a solution contains more than eighty 
times as many OH~ as H + ions it colors phenolphthalein red 
(the alkaline color) ; if it contains less than eighty times as many 
OH~ ions as H + ions it leaves phenolphthalein colorless. On 
the other hand, methyl orange shows an orange color (its inter- 
mediate shade between pink, the acid color, and yellow, the 
alkaline color) when the number of H + ions is about a million 
times the number of OH~ ions. Anomalous as it may seem at 
first thought, it is really fortunate that many of our indicators 
do not indicate perfect neutrality; for suppose we wish to dis- 
cover how much acetic acid a certain solution contains. We 
may titrate it accurately with normal or decinormal sodium 
hydroxide or other strong base if we use the right indicator 
(137) . Now we have learned that when acetic acid is mixed with 
exactly the theoretically equivalent amount of NaOH the result- 
ing solution is not perfectly neutral but in reality slightly alkaline 



272 



Introduction to General Chemistry 



(434). In accord with this we found that a solution of NaC 2 H 3 2 
was slightly alkaline to litmus, showing that the number of OH~ 
ions was greater than the number of H + ions. Therefore we 
must use as a titration indicator one which shows its change of 
color when the number of OH~ is greater than the number of H + 
ions. We find that phenolphthalein proves to be just right for 
the purpose. In general, we use phenolphthalein as indicator 
in titrating all moderately weak acids. 

If we wish to titrate NH 4 OH with HC1 we cannot use 
phenolphthalein, because a solution of NH 4 C1 contains more H + 
than OH~ ions. Such a solution " seems" acid to this indicator. 
We must use one which changes color when the number of H + 
ions exceeds the number of OH~ ions, and for this case we find 
methyl orange satisfactory. In general, we use methyl orange 
in titrating moderately weak bases. The acid used in such 
titrations must be a strong one. In titrating a strong base with 
a strong acid any of these indicators gives sufficiently accurate 
results. Table XVIII gives the colors of indicators in solutions 

TABLE XVIII 



Hydrogen ion concentration 
Hydroxyl ion concentration. 

Metrryl orange 

Litmus 

Phenolphthalein 



IO — 3 


IO — s 


10-7 


10- 11 

Pink 
Red 
Colorless 


10-9 
Yellow 
Red 
Colorless 


10-7 
Yellow 
Purple 
Colorless 



10- 8 

10-6 

Yellow 

Blue 

Red 



of hydrogen and hydroxyl ion concentrations near those at 
which the color change occurs. In this table the concentrations 
are given in gram molecular weights per liter. If the H + con- 
centration is io -3 , 1,000 liters contain 1 g. of H + ion. 

441. Summary on Equilibrium between Soluble Electro- 
lytes. — If we mix solutions of two electrolytes, AB and CD, 
having no ion in common, a double decomposition reaction, 

AB+CD±zAD+CB, 



takes place to a greater or less extent, because of the tendency 
of each positive ion to combine with each negative ion present. 



Applications of the Ionic Hypothesis 273 

If all four substances of the preceding equation are highly 
ionized (Class I, 415), the mixed solution will contain largely 
the four sorts of free ions, A + , B~ , C + , andZ>~. Only a small 
percentage of the dissolved material will be present as molecules. 
If all four substances are equally ionized, equal numbers of 
molecules of each will be present, as shown in Fig. 56. 

If one of the four substances (say AD) is little ionized (Class 
II, 419), then the large numbers of A + and D~ ions shown in 
Fig. 56 cannot exist side by side in the mixed solution, since they 
will very largely combine to form AD molecules. The disap- 
pearance of A + and D~ ions allows AB and CD molecules more 
or less completely to dissociate. The final result, shown in 
Figs. 59, 61, and 63, is a nearly complete reaction, AD being 
present almost wholly in un-ionized form and CD to a small 
extent as molecules, but largely as C + and D~ ions. 

A generalization of much importance is found in the Common 
Ion Law: suppression of the ionization of a little ionize sub- 
stance occurs when we add to its solution a highly ionized 
substance having a common ion. 

Since* in pure water the number of H + ions is equal to the 
number of OH~ ions, and since we may consider pure water a 
perfectly neutral substance, we define a neutral solution as one 
in which the number of H + ions is exactly equal to the number 
of OH~ ions. 

Class III (434, 437) comprises reactions in which one little 
ionized substance reacts with a highly ionized substance to 
form products one of which is slightly, the other highly, ionized. 
Examples are found in the neutralization of a weak acid by a 
strong base; or of a weak base by a strong acid. In such cases 
the reaction is more or less incomplete. The weaker the acid 
or base taken, the less complete is the neutralization. Con- 
versely, salts of weak acids or of weak bases are hydrolyzed by 
water. The former give solutions which are alkaline, the latter 
those which are acid, in reaction. This kind of action is called 
hydrolytic dissociation. 

Under Class IV (438) it was pointed out that weak acids and 
weak bases always react incompletely, and that when either 



274 Introduction to General Chemistry 

or both are extremely weak, salt formation may not occur in 
solution (177). 

We have seen that indicators change color according to the 
concentration of H + and OH - ions present. Litmus shows 
its neutral tint when the numbers of H + and OH~ ions are 
nearly equal. Phenolphthalein requires an excess of OH~ 
ions to change color, while methyl orange requires an excess of 
H + ions. 



CHAPTER XIX 

APPLICATIONS OF THE IONIC HYPOTHESIS. REACTIONS 
INVOLVING CHANGES OF STATE 

442. Introduction. — In the present chapter we shall study 
precipitation from the standpoint of the ionic hypothesis in 
order to understand the underlying principle's of this most 
important means of separating substances. In equations for 
precipitation reactions, the substance precipitated will be indi- 
cated by a downward-pointing arrow. 

If we consider the familiar examples of precipitation repre- 
sented by the following equations, 

AgN0 3 +HCl^AgCty+HN0 3 , 

BaCl 2 +H 2 S0 4 ^BaS0 4 |-f 2 HCl, 

we might conclude that AgCl and BaS0 4 are precipitated because 
they are insoluble in water. We might even be tempted to say 
that in the reaction 

AB+CD^AD+CB, 

if either AD or CB is an insoluble substance it will be pre- 
cipitated. This statement contains something of the truth, 
but it is far from the whole truth, as the following examples will 
prove. Calcium carbonate, CaC0 3 (marble), is an almost 
insoluble substance. If we mix solutions of calcium chloride 
and carbonic acid we might expect to get a precipitate of calcium 
carbonate, thus, 

CaCl 2 +H 2 C0 3 ^CaC0 3 ^+2HCl. 

Not a trace of precipitate is formed. On the other hand potas- 
sium chlorate, KC10 3 , is easily soluble in water; but if we add 
a saturated solution of potassium bromide, KBr, to a saturated 
solution of sodium chlorate, NaC10 3 , a precipitate of KC10 3 
forms. Evidently the matter is not as simple as at first thought 
it appears to be. The separation of a solid from a solution is 

275 



276 Introduction to General Chemistry 

obviously the reverse of the passage of a solid into solution. 
Accordingly, in beginning the study of precipitation, it will be 
advisable for the student to read again sections 120-23. In 
section 122 it is stated, a A solution which at a fixed tempera- 
ture will dissolve no more of a given substance is called a 
saturated solution. When we speak of the solubility of a sub- 
stance we mean the amount- of substance dissolved in a given 
amount of water in the case of the saturated solution." 

443. The Kinetic Theory of Solution. — When a soluble salt 
is brought into water its molecules begin to leave the surface 
of the solid and pass into the water. Immediately thereafter 
dissolved salt molecules will occasionally strike the surface 
of the solid and in some cases remain attached thereto. Finally, 
when the solution has become saturated we may imagine that 
the equilibrium between dissolved and solid salt is the result 
of the passage of molecules into and out of solution at exactly 

equal rates, thus: 

' AB±*AB 
Solid Dissolved 

This picture is, however, incomplete, since the salt is partly 
ionized. The dissolved molecules are therefore in equilibrium 
with their ions as well as with the solid salt, thus : 

AB^AB^A++B~ 

Solid Dissolved 

444. Graphic Representation of a Solid Electrolyte in 
Equilibrium with Its Saturated Solution. — We shall represent 

NaCI 





Fig. 70 

a solid electrolyte (acid, base, or salt) graphically by a cross- 
hatched square. The condition of a saturated solution of a 
soluble salt (NaCI, for example) in contact with an excess of 
the solid salt may then be represented as in Fig. 70. 



Applications of the Ionic Hypothesis 277 

445. The Solubility of Molecules. Molecular Solubility.— 
If we except a small number of electrolytes like sulfuric and 
nitric acids, which mix with water in all proportions, all other 
acids, bases, and salts have limited solubilities in water. Since 
all electrolytes are more or less ionized in solution, the dissolved 
substance is present partly as molecules and partly as ions. 
Therefore the total solubility of a substance in a solution saturated 
at a given temperature is the sum of the solubility of its molecules 
and the solubility of its ions. It seems reasonable to assume 
that the limited solubility of an electrolyte as a whole is the 
result of the limited solubility of its molecules rather than of its 
ions. Two reasons may be given for this assumption which will 
be amply confirmed by additional evidence to be considered 
later. 

In the first place the solid salt passes into and out of solution 
as molecules (see Fig. 70). If the molecules have a limited 
solubility, this would limit the solubility of the ions as well, since 
the latter and the former are directly in equilibrium. Therefore 
it is sufficient to assume limited solubility of the molecules in order 
to explain limited total solubility. Secondly, the small solubility 
of a difficultly soluble salt like CaS0 4 (100 c.c. of water dissolve 
o.25g. of CaS0 4 ) cannot be due to a correspondingly small 
solubility of Ca +H ~ or S0 4 ions, since solutions of CaCL and 
many other easily soluble and highly ionized calcium salts con- 
tain an abundance of Ca ++ ions, and solutions of H 2 S0 4 and 
many easily soluble and highly ionized sulfates contain large 
concentrations of S0 4 ~~ ions. We shall assume, therefore, that 
at a given temperature the solubility of an acid, base, or salt is 
limited by the solubility of its molecules; and we shall call the 
solubility of the molecules (in the saturated solution) the 
molecular solubility (abbreviated M.S.) of the substance. Sum- 
marizing, we may say that when a solid electrolyte is mixed with 
water at a fixed temperature the substance dissolves and the 
concentration of the solution increases until the M.S. is reached; 
the solution is then saturated (at that temperature), and the 
molecules are in equilibrium with the ions and with the solid 
substance. 



278 



Introduction to General Chemistry 



- - v CD 

\ 




H 



CB 



AD 



446. The Cause of Precipitation. — We are now ready to 
apply the foregoing principles, together with those learned in 
chapter xviii, to the process of precipitation. We have learned 
(414) that in the reaction 

AB+CD^AD+CB, 

if all four substances are easily soluble and highly ionized the 
resulting solution contains largely the four sorts of ions, together 

with small proportions of the 
four kinds of molecules. The 
conditions before and after 
the reaction are shown in Figs. 
55 and 56. Now let us suppose 
that one of the products AD 
is not very soluble, so that its 
molecular solubility (M.S.) is 
less than that corresponding to 
the concentration of the mole- 
cules of AD formed in reaction 
(Fig. 56). In this case the 
AD amount of molecules of AD 
formed in excess of the M.S. 
will separate out of solution as 
a precipitate. Fig. 71 shows 
the resulting condition for the case where the M.S. of AD is 
rather small but not extremely small. By comparison of Figs. 
56 and 71 we see that an appreciable shift in equilibrium of the 
dissolved substances accompanies the partial precipitation of AD. 

447. The Precipitation of KC10 3 . — An actual example con- 
forming perfectly to the conditions set forth in the preceding, 
paragraph is found in the reaction 

KBr+NaC10 3 ±5KC10 3 |+NaBr. 

Of the four salts, all are very soluble except KC10 3 , which dis- 
solves only to the extent of 7 g. in 100 c.c. of water at 18 . All 
four salts are highly and about equally ionized in solutions of 
equal concentration. The conditions of the solutions of KBr 



n 



Fig. 71 



Applications of the Ionic Hypothesis 279 

and NaC10 3 before mixing are shown with sufficient accuracy 
by Fig. 55, while Fig. 56 shows the condition which the mixed 
solution would reach if KC10 3 were also very soluble. It hap- 
pens, however, that the amount of molecular KC10 3 which 
tends to be formed exceeds the rather small M.S. of this sub- 
stance, and in consequence the excess above the M.S. separates 
as a precipitate. Precipitation continues until the amount of 
molecular KC10 3 left in solution is equal to the M.S. of this sub- 
stance. The mixture is then in the condition of equilibrium 
shown in Fig. 71. Comparison of Figs. 56 and 71 shows that 
the removal (by precipitation) of KC10 3 from the solution causes 
a marked shift in the equilibrium. We may trace the stages as 
follows: Fig. 56 shows the condition that would exist if no 
precipitation occurred. The removal of KC10 3 , results in the 
further union of K + and C10 3 ~ ions to form more KC10 3 . The 
resulting loss of K + and C10 3 ~ ions promotes the further ioniza- 
tion of KBr and NaC10 3 respectively and thus increases the 
numbers of Br~ and Na + ions. The latter ions unite in part 
to form additional molecular NaBr. The final result is the 
change from the condition of Fig. 56 to that of Fig. 71. The 
principles here exemplified apply to all double decomposition 
precipitations. 

448. The Precipitation of CaC0 3 . —Let us now consider a 
case in which one of the products is precipitated almost com- 
pletely. The reaction 

CaCl 2 +Na 2 C0 3 ^CaC0 3 i+2NaCl, 

in which CaC0 3 is the precipitate, will serve as a typical illustra- 
tion. Although CaC0 3 appears to be insoluble in water, it is in 
fact slightly soluble, and has therefore a definite but very small 
M.S. The other three salts, CaCl 2 , Na 2 C0 3 , and NaCl, are 
easily soluble and highly ionized, and in consequence the reaction 
between solutions of CaCl 2 and Na 2 C0 3 tends to reach the con- 
dition shown in Fig. 56 illustrating a Class I reaction (414). In 
this respect it completely resembles the reaction 

KBr+NaC10 3 ==>KC10 3 |+NaBr. 



280 



Introduction to General Chemistry 



- Na=CO, 




It differs however from this reaction in that the M.S. of CaC0 3 
is extremely small compared with the M.S. of KC10 3 . In con- 
sequence the CaC0 3 formed pre- 
\ co " cipitates almost completely, as 
7 illustrated in Fig. 72. In all 
double decomposition reactions 
~-NCa++ of the above-mentioned types 
1 (all involved substances highly 
"-' ionized) the precipitation is the 

n more complete the less the M.S. 

' ^ ^cacoa of the precipitate . 
- 1 449. The Action of H 2 C0 3 

on CaCl 2 . — In section 442 it 
u was pointed out that H 2 C0 3 

CaC03 does not precipitate CaCl 2 , as 
we might expect according to 
the following hypothetical 
Fig. 72 equation: 



n 



^ v NaCI 




CaCl 2 +H 2 C0 3 ^CaC0 3 ^+2HCl. 



The reason is as follows: carbonic acid H 2 C0 3 is a very weak 
acid and in consequence yields but very few C0 3 ~~ ions; and 
although CaCl 2 gives an abundance of Ca ++ ions, the concen- 
tration of 003"" ~ ions is so small that the concentration of 
CaC0 3 molecules formed is less than the M.S. of this substance. 
Therefore no precipitation of CaC0 3 takes place. The differ- 
ence in behavior of H 2 C0 3 and Na 2 C0 3 toward a solution of 
CaCl 2 is wholly due to the difference in their tendencies to ionize, 
in consequence of which a solution of H 2 C0 3 contains exceedingly 
few C0 3 ~~ ions as compared with a solution of Na 2 C0 3 . 

The behavior of H 2 C0 3 is typical of that of all weak (little 
ionized) electrolytes. In the precipitation of salts, weak acids 
and bases are, in general, less efficient precipitants than their 
salts, since the latter are highly ionized. 

450. The Precipitation of Magnesium Hydroxide, Mg(OH) 2 . 
— We shall next discuss in detail the precipitation of Mg(OH) 2 , 



Applications of the Ionic Hypothesis 281 

not so much because of the chemical importance of this sub- 
stance, but because the reactions illustrate in a striking way some 
of the most important principles of ionic equilibrium. 

If we add NaOH to a solution of magnesium chloride, 
MgCl 2 , we obtain an abundant white precipitate of Mg(OH) 2 , 
formed as follows: 

MgCl 2 +2NaOH^Mg(OH) 2 |+ 2 NaCl. 

If we use NH4OH instead of NaOH the reaction is similar but 
reaches a state of equilibrium when only a part of the magnesium 
is precipitated : 

MgCl 2 + 2NH 4 OH±? Mg(OH) 2 |+ 2 NH 4 C1. 

If we add to a MgCl 2 solution a solution of NH 4 OH mixed with 
sufficient NH 4 C1, no precipitation occurs. We shall now explain 
these facts. In the first place we may say that the action of 
NaOH on MgCl 2 is analogous to the action of Na 2 C0 3 on CaCl 2 , 
the two equations being 

MgCl 2 + 2Na0H±5 Mg(0H) 2 sH- 2 NaCl, 
CaCl 2 + Na 2 C0 3 s* CaC0 3 |+ 2 NaCl. 

Both MgCl 2 and NaOH, like CaCl 2 and Na 2 C0 3 , are easily solu- 
ble and highly ionized; sodium chloride, one of the products in 
both reactions, is also highly ionized. The other product in the 
first reaction, Mg(0H) 2 , is but slightly soluble, and like CaC0 3 
it is therefore precipitated almost completely. 

If, however, we use NH 4 OH instead of NaOH, the reaction is 
far from complete. The reason can best be seen by the aid of 
Figs. 73 and 74. Figure 73 shows the condition of the solutions 
before they are mixed; Fig. 74 represents the condition of the 
mixture. 

It will be recalled, as. shown in Fig. 73, that NH 4 OH is but 
little ionized. Still its solution yields sufficient OH~ ions to 
produce in reaction with magnesium chloride solution more 
Mg(0H) 2 than the M.S. of the latter difficultly soluble substance. 
The excess of Mg(OH) 2 above its M.S. precipitates, Fig. 74. As 
these changes go on, molecules of NH 4 OH continue to ionize, thus 



282 



Introduction to General Chemistry 



bringing into the solution far more NH 4 + ions than were 
originally present (cf. Figs. 73 and 74). The presence of the 
large excess of NH 4 + ions restricts the number of OH~ ions to 
such an extent that a state of equilibrium is reached in reaction, 
Mg(OH) 2 (dissolved)^Mg+++20H- 

while there is still a considerable amount of magnesium in the 
form of Mg ++ ions and MgCl 2 molecules left in the solution. 
After this condition is reached 
no more Mg(0H) 2 precipi- 
tates. We therefore conclude 
that NH 4 OH precipitates 
Mg(0H) 2 only partially, (1) 
because NH 4 OH is a weak or 
little ionized base, and (2) 




MgCI 2 



NH.OH 




n 



NH, 4 



NHXI 



u* 



Mg(OH), 





± I 

\ / 

V / 



/ ~~ 


\ 

/ 


/ 
\ 




~ - S 0H" 
\ 

/ 



1 Mg(OH), 



iiiiiiiiiiiiiiiniimi 



Fig. 73 



Fig. 74 



because the accumulation of NH 4 + ions (Fig. 74) finally restricts 
the OH~ concentration to so small a value that the Mg ++ 
and OH~ ions are just in equilibrium with the amount of 
Mg(OH) 2 corresponding to its M.S. It will now be easy to 
understand why no Mg(0H) 2 is precipitated when NH 4 OH, 
mixed with considerable NH 4 C1, is added to a MgCl 2 solution. 

The NH 4 C1 furnishes at once such an excess of NH 4 + ions 
that the OH~ concentration is decreased to so small a value that 
less Mg(0H) 2 is formed in the reaction 

Mg+++20H"±>Mg(OH) 2 

than corresponds to its M.S. Therefore no precipitation occurs. 

451. The Precipitation of Ferric Hydroxide, Fe(OH) 3 . — The 

reaction 

FeCl 3 +3NH 4 OH^Fe(OH) 3 ^+NH 4 Cl 



Applications of the Ionic Hypothesis 283 

takes place with the practically complete precipitation of brown 
Fe(OH) 3 , which is almost insoluble in water. The completeness 
of precipitation is not noticeably affected by the addition of much 
NH 4 C1. There are two reasons for the difference in behavior of 
Fe(OH) 3 and Mg(OH) 2 : (1) the former is much more insoluble 
in water than the latter, so that the M.S. of the latter (although 
small) is perhaps 1,000 times as large as that of the former; 
(2) Mg(0H) 2 is a rather strong (highly ionized) base, while 
Fe(OH) 3 is a very weak base. Therefore in the reactions 

Mg+++ 2 OH-^Mg(OH) 2 (dissolved), 
Fe++++ 3 OH-±5Fe(OH) 3 (dissolved), 

for equal concentrations of Mg ++ , Fe +++ , and OH~ ions far 
less Mg(0H) 2 is formed than Fe(OH) 3 . The presence of NH 4 C1 
decreases the OH~ concentration of NH 4 OH, but not sufficiently 
to prevent the practically complete precipitation of Fe(OH) 3 
because of the weakness of the latter and its exceedingly small 
M.S. The weaker a base and the smaller its M.S., the more com- 
pletely is it precipitated by N H 4 OH, and the less its precipitation 
is hindered by the presence of ammonium salts. 

452. Classification of Precipitations. — The various examples 
of precipitation just studied cover the important principles con- 
cerned. We may now cite a few additional examples of each of 
these classes of precipitation. In the reaction 

KBr+NaC10 3 ±5KC10 3 |+NaBr, (447) 

all four substances are highly ionized, and all but KC10 3 are very 
soluble. The latter is partially precipitated because it is formed 
in excess of its not very large M.S. The following reactions are 
of this class: 

CaCl 2 +2NaC10 3 ±52NaCl|+Ca(C10 3 ) 2 , 
P.b(N0 3 ) a +2NaCl±5PbCU+2NaN0 3 , (167) 

CaCl 2 +Na 2 S0 4 ±sCaS0 4 |+2NaCl. (153) 

In the first reaction saturated solutions are required to give a 
precipitate of NaCl. 



284 Introduction to General Chemistry 

In the second example studied, 

CaCl 2 +Na 2 C0 3 ^CaC0 3 ^+2NaCl, (448) 

the precipitate CaC0 3 has an extremely small M.S. Its pre- 
cipitation by Na 2 C0 3 is almost complete. Other reactions of 
this class are: 

AgN0 3 +NaCl = AgCty+NaN0 3 , (382) 

AgN0 3 +KBr = AgBr|+KN0 3 , (257) 

Pb(N0 3 ) 2 +CuS0 4 = PbS0 4 |+Cu(N0 3 ) 2 , (167) 

BaCl 2 +Na 2 S0 4 = BaS0 4 |+ 2NaCl, (164) 

MgCL+ 2 NaOH = Mg(OH) 2 |+ 2 NaCl. ' (450) 

In the third example (449) it was shown that H 2 C0 3 did not 
give with CaCl 2 a precipitate of CaC0 3 , because the former is 
very little ionized and therefore yields very few G0 3 ions. The 
following pairs of substances also fail to give precipitates, because 
in each case of the weakness of the acid coupled with the moderate 
solubility of the salt, that might be precipitated : 

CaCl 2 and H 3 P0 4 , 
FeCl 2 and H 2 S, " 
AgN0 3 and HC 2 H 3 2 . 

On the other hand the reactions 

3 CaCl 2 +2Na 3 P0 4 = Ca 3 (P0 4 ) 2 ^+6NaCl, (158) 

FeCl 2 + (NH 4 ) 2 S = FeS^+ 2 NH 4 C1, 
AgN0 3 +NaC 2 H 3 2 = AgC 2 H 3 2 |+NaN0 3 , 

give abundant precipitates, because, instead of the weak acids, 
we use their salts, which are highly ionized. 

The fourth example, which belongs to Class II, was taken up 
in contrast to the fifth example, which is typical of a fourth class 
of precipitation reactions. The fourth and fifth examples were: 

MgCl 2 +2NaOH^Mg(OH) 2 sk+2NaCl, (450) 

MgCl 2 +2NH 4 OH^Mg(OH) 2 |+2NH 4 Cl. (450) 

The precipitation of Mg(0H) 2 is nearly complete in the first 
reaction but only partial in the second, owing to the moderate 
solubility of Mg(OH) 2 and the little ionization of NH 4 OH, 
especially in the presence of its salts. 



Applications of the Ionic Hypothesis 285 

The following reaction of a weak electrolyte (H 2 S) also results 
in partial precipitation: 

ZnCl 2 +H 2 S ^ZnS|+ 2HCI. 

In this case the precipitation is prevented by the presence of 
an excess of HC1 or other strong acid, because of the suppression 
of the ionization of the H 2 S by the H + ions of the strong acid. 
The sixth example dealt with the precipitation of Fe(OH) 3 : 

FeCl 3 + 3 NH 4 OH^Fe(OH) 3 |+ 3 NH 4 Cl. (451) 

In this case the precipitate is so insoluble (M.S. so small) and 
so weak (little ionized) that it is practically completely pre- 
cipitated by NH 4 OH even in the presence of NH 4 C1. Although 
we call NH 4 OH a weak base, it is enormously stronger than 
Fe(OH) 3 , even when mixed with much NH 4 C1. Other reactions 
which fall into this class are: 

A1C1 3 + 3 NH 4 0H^A1(0H) 3 |+ 3 NH 4 C1, (174) 

CrCl 3 + 3 NH 4 OH^Cr(OH) 3 ^+ 3 NH 4 Cl, (344) 

CuS0 4 +H 2 S^CuS|+H 2 S0 4 , 

Ag 2 S0 4 +H 2 S^Ag 2 S|+H 2 S0 4 . 

Excess of NH 4 C1 in the first two cases, and of HC1 or H 2 S0 4 in 
the last two cases, fails to prevent practically complete precipita- 
tion. 

453. Precipitation by Adding a Substance Having a Common 
Ion. — We have learned in the foregoing chapter (432) that if we 
add to the solution of an electrolyte, AB, enough of another 
highly ionized electrolyte, AC, having a common ion, A, to 
increase the concentration of the common ion, the degree of 
ionization of the first substance will be suppressed. If now the 
substance A B is not very soluble, the suppression of its ioniza- 
tion caused by adding AC may increase the concentration of 
the A B molecules to such an extent as to exceed the M.S. of AB. 
In consequence part of AB will separate out as a precipitate. 
For example, if a few bubbles of HC1 gas are passed into a 
saturated solution of NaCl, a precipitate of NaCl is formed. A 
similar result is also produced by adding a little concentrated 
HC1 to a saturated salt solution. In each case the ionization 



286 Introduction to General Chemistry 

of the salt is suppressed by reason of the increase in concentra- 
tion of the Cl~ ions, and the concentration of the molecular NaCl 
increased beyond the M.S. of this substance. Salt precipi- 
tates until the concentration of molecular NaCl falls to the 
value corresponding to its M.S. Another example illustrating 
the same principle is found in the precipitation of KC10 3 from 
its saturated solution by the addition of a saturated solution 
of either KBr or NaC10 3 . We may say that as a general rule the 
total solubility of a salt (molecular and ionic) is diminished by the 
presence in the solution of another electrolyte having a common ion. 
454. Conditions Favoring Precipitation. — In the reaction 

AB+CD=AD+CB 

precipitation will occur if one of the products, say AD, is formed 
as molecules in greater concentration than its molecular solu- 
bility. In brief, if the M.S. is exceeded, precipitation will occur. 
Now the smaller the M.S. of AD, the more probably will it be 
exceeded. 

On the other hand the M.S. is the more likely to be exceeded 
the greater the concentration of AD which tends to be produced 
in the reaction. The various factors which determine the 
amount of AD produced (when AD is soluble) have been dis- 
cussed at length in chapter xviii (Summary, 441). 

These applications of the ionic hypothesis have the following 
bearing on the practice of precipitation. In the first place, if 
we are to precipitate from solution an "insoluble" salt of a weak 
acid we use as the precipitant a soluble salt of the weak acid 
instead of the acid itself, since the former is highly ionized, 
while the latter is not. (The term precipitant means the reagent 
added to cause precipitation.) If, however, we are to precipitate 
an insoluble chloride we may use either a soluble chloride or 
hydrochloric acid, since this strong acid is as highly ionized as 
its salt. When the precipitant is added to a given solution, a 
precipitate may not appear until considerable reagent has been 
added. When it is no longer possible to see if more precipitate 
is forming with further additions of the reagent, a small portion 
of the mixture is filtered, or, better, the precipitate is allowed 



Applications of the Ionic Hypothesis 287 

to settle, and the clear solution is tested with more of the pre- 
cipitant. Only moderate excesses of the precipitant are used 
as a rule, since in many cases the precipitant reacts farther with 
the precipitate to form new and soluble compounds, with the 
result that the precipitate dissolves in an excess of the reagent 
added. 

455. Dissolving Solid Substances. — Substances which are 
not readily soluble in water often dissolve easily in solutions of 
other electrolytes. In such cases we may imagine that chemical 
reaction gives rise to new products which are soluble in water. 
Here is a case in point: Calcium hydroxide, Ca(0H) 2 , is but 
slightly soluble in water (o.i2g. in iooc.c), giving a very 
dilute solution known as limewater. If we mix a few grams of 
Ca(OH) 2 with 100 c.c. of water, most of the solid remains undis- 
solved. If now we add dilute HC1 to the mixture, the solid 
finally passes completely into solution. The explanation is as 
follows: The small amount of dissolved Ca(0H) 2 (which is a 
strong base) is neutralized by the added HC1 to form very soluble 

Ca(OH) 2 + 2 HC1±5 CaCl 2 + 2 H 2 0. 

More Ca(OH) 2 then dissolves in the water in the tendency to 
keep the concentration of the dissolved Ca(0H) 2 up to its M.S.: 

Ca(OH) 2 ^ Ca(OH) 2 ±5 Ca+++ 2 OH" 
Solid Dissolved 

As fast as Ca(0H) 2 passes into solution it reacts with the HC1 
present. If the chemically equivalent amount of HC1 is added, 
all Ca(0H) 2 will finally dissolve, and the solution will consist 
simply of CaCl 2 dissolved in water. 

A perfectly analogous reaction is found in the dissolving of 
the difficultly soluble, strong base Mg(0H) 2 in dilute HC1: 

Mg(OH) 2 + 2 HC1^ MgCl 2 + 2 H 2 0. 

Even if the base is weak and much less soluble than either 
Ca(OH) 2 or Mg(0H) 2 , it will usually dissolve in water upon the 
addition of a strong acid. For example, Fe(OH) 3 is a very weak 



2&8 Introduction to General Chemistry 

base almost insoluble in water; it dissolves readily in dilute 
HC1, forming a solution of ferric chloride, 

Fe(OH) 3 + 3 HCl±5FeCl 3 +3H 2 0. 

The stages in the process of dissolving may be considered to be 
comparable to those in the case of the dissolving of Ca(0H) 2 in 
dilute HC1. 

Most bases, with the exception of the hydroxides of sodium, 
potassium, ammonium, and barium, are very little soluble 
in water. All such so-called insoluble bases dissolve in dilute 
HC1, HN0 3 , an d H 2 S0 4 to form clear solutions, if their cor- 
responding salts with these acids are soluble in water. 

456. Dissolving Little Soluble Salts of Weak Acids by Solu- 
tions of Strong Acids. — Silver acetate is a rather difficultly solu- 
ble salt (i.og. dissolves in 100 c.c. H 2 at 18 ) which is easily 
made by precipitating AgN0 3 with NaC 2 H 3 2 , 

AgN0 3 +NaC 2 H 3 2 ^AgC 2 H 3 2 +NaN0 3 . (452) 

If we mix 3 or 4 g. of AgC 2 H 3 2 with 100 c.c. of H 2 0, only a small 
portion dissolves; but upon addition of dilute HX0 3 the whole 
of the solid passes into solution. Silver acetate is the salt of 
the weak acid HC 2 H 3 2 , and, as we have already learned (428), 
a strong acid reacts more or less completely with the (soluble) 
salt of a weak acid to form the weak acid and the salt of the strong 
acid. This was shown earlier in the case of the reaction 

HCl+NaC 2 H 3 2 ±5HC 2 H 3 2 +NaCl. (424) 

Xitric acid reacts similarly with the dissolved portion of the 
AgC 2 H 3 2 , 

HN0 3 +AgC 2 H 3 2 ^HC 2 H 3 2 +AgN0 3 . 

The reaction is nearly complete, and both products are easily 
soluble. The dissolved molecular AgC 2 H 3 2 being thus removed 
from the solution, more of the solid passes into solution in the 
tendency to keep the concentration of molecular AgC 2 H 3 2 up 
to its M.S. But as this salt reacts with the HX0 3 present as 



Applications of the Ionic Hypothesis 289 

soon as it comes into solution, its M.S. is never reached, so that 
finally all of the solid passes into solution. The solution con- 
sists largely of AgN0 3 and its ions, together with molecular 
acetic acid. 

In many other cases so-called "insoluble" salts of weak acids 
dissolve in solutions of strong acids like HC1, HN0 3 , and H 2 S0 4 . 
The following reactions are of this type: 

FeS+2HCl = H 2 S+FeCl 2 , 
CaC0 3 + 2HCI = H 2 C0 3 + CaCl 2 , 

Ca 3 (P0 4 ) 2 +6HCl = 2H 3 P0 4 + 3 CaCl 2 . 

However, not all "insoluble" salts of weak acids dissolve in 
strong acids. For example, CuS, which comes down as a black 
precipitate when H 2 S is passed into a solution of a copper salt, 
and is therefore a salt of the very weak acid H 2 S, does not dissolve 
appreciably in cold HCL The reason for this is directly trace- 
able to the extremely small M.S. of CuS. In general the smaller 
the M.S. of a salt of a weak acid the less soluble it is in a strong 
acid. Other examples of this sort are found in Ag 2 S and HgS, 
neither of which is dissolved appreciably by dilute HC1 or 
H 2 S0 4 . 

457. Weak Acids and Salts of Strong Acids. — We have 
already learned (282) that the equilibrium mixture has the same 
composition whether we start with one pair of substances of a 
reaction or the equivalent amounts of the other pair. In accord 
with this principle we always find that if a reaction takes place 
practically completely in one direction, the reverse of the reaction 
does not succeed under the same conditions of temperature and 
concentration. In sections 449 and 452 it was stated that 
mixtures of the following pairs of substances fail to give pre- 
cipitates, although little soluble salts would be formed by 
double decomposition : 

H 2 C0 3 and CaCl 2 , 
H 3 P0 4 and CaCL, 
H 2 S and FeCL, 
HC,H 3 2 and AgN0 3 . 



290 Introduction to General Chemistry 

Therefore we may be certain that calcium carbonate, calcium 
phosphate, and ferrous sulphide are soluble in hydrochloric acid, 
and that silver acetate is soluble in nitric acid. Since the deter- 
mining factor in dissolving each of these salts is the formation 
of the weak acid, we may go farther and predict that any strong 
acid will dissolve these salts. Sometimes a new insoluble salt 
is formed by the strong acid, as when hydrochloric acid acts on 
silver acetate; but such reactions are secondary to the solution 
of the original salts. 

458. "Insoluble" Salts of Strong Acids.— The "insoluble" 
salts of strong acids are not as a rule dissolved to an appre- 
ciable extent by solutions of other strong acids. For example, 
AgCl is not appreciably dissolved by HN0 3 , although the 
products HC1 and AgN0 3 of the hypothetical reaction 

AgCl+HN0 3 ^HCl+AgN0 3 

are both easily soluble substances. A reaction in which both 
of the products are highly ionized, as in this case, falls in Class I 
(414). In all such reactions very little chemical change occurs, 
and this is more strikingly true the more dilute the solution. As 
we are now considering the case where one of the substances 
taken is nearly insoluble in water, the solution of this substance 
must be exceedingly dilute. Comparing the action of HN0 3 
on AgC 2 H 3 2 and AgCl, we may say that the first reaction takes 
place readily because of the tendency of H + and C 2 H 3 2 ~ to 
unite nearly completely to form little ionized HC 2 H 3 2 ; and 
that the second reaction does not progress far because of the 
slight tendency for H + and Cl~ ions to unite, since HC1 is 
nearly completely ionized in very dilute solution. 

459. Evolution of a Gas. — Substances may separate from 
solutions in two ways: (1) as solid precipitates and (2) as gases. 
We have considered the first case and shall now take up the 
second, and we shall see that if a substance separates from a 
solution as a gas the effect on the ionic equilibrium is the same 
as if the substance separated as a solid. The principles that 
apply to precipitation apply also, with slight obvious modifica- 
tions, to gas evolution. Gases have limited solubilities, and 



Applications of the Ionic Hypothesis 291 

instead of the M.S. of the precipitate we have the molecular 
solubility (M.S.) of the gas. Let us now consider a few well- 
known reactions as illustrations. 

460. The Action of H 2 S0 4 on NaCl. — If we mix dilute solu- 
tions of H 2 S0 4 and NaCl no visible effect is produced, although 
in the solution the reaction 

H 2 S0 4 +NaCl^HCl+NaHS0 4 

takes place partially. This is a Class I reaction (414) since all 
four substances are easily soluble and highly ionized. Therefore 
the dilute solution contains chiefly the ions and relatively few 
molecules. Nevertheless some HC1 molecules are formed even 
in dilute solution, but, as HC1 is a very soluble gas, none of it 
escapes from the solution. 

On the other hand the results are quite different if but little 
water is present. In making hydrochloric acid (103) 58 g. of 
NaCl, 100 g. of concentrated H 2 S0 4 , and 30 g. of water were 
mixed in a flask and gently heated. 

The proportions of NaCl and H 2 S0 4 taken were those indi- 
cated by the foregoing equation, since the 100 g. of concentrated 
acid taken consists of 98 g. of actual H 2 S0 4 and 2 g. of H 2 0. 
If the reaction should take place completely, 36. 5 g. of. HC1 
would be formed. This is far more HC1 than the 32 g. of water 
present can hold in solution, especially when the mixture is 
heated. Therefore HC1 gas escapes from the solution. ' The loss 
of HC1 by the solution impedes the reverse action on the 
NaHS0 4 present and so causes a great shift to the right of the 
equilibrium that would otherwise be reached in the reaction 

H 2 S0 4 +NaCl^HCli+NaHS0 4 . 

As a consequence this reaction goes nearly completely from left 
to right under the conditions described (103), the HC1 being 
given off as gas. In equations for reactions involving gas 
evolution the gas will be indicated by an upward-pointing arrow. 

461. The Action of HC1 on CaCO,. — We have already seen 
that carbonic acid, H 2 C0 3 , does not precipitate CaC0 3 from a 
CaCl 2 solution, and have learned that this is because H.CO- is 



292 Introduction to General Chemistry 

so little ionized that insufficient molecular CaC0 3 is formed to 
exceed its M.S. This fact indicates that the reactions 

2HCI+ CaC0 3 ^ CaCl 2 +H 2 C0 3 , 
H 2 C0 3 ^H 2 0+C0 3 f, 

will take place practically completely, since in all reactions 
between electrolytes exactly the same proportions of the same 
products result, whether we start with one pair of substances 
or the chemically equivalent amounts of the other pair (282). 
The dissolving of CaC0 3 in dilute HC1 takes place as follows: 
CaC0 3 first dissolves to the limit of its (very small) M.S. in the 
water present; the dissolved molecules then ionize rather highly: 

CaC0 3 ±5Ca+++C0 3 "-. 

The C0 3 ions unite nearly completely with the H + ions of 
the highly ionized HC1 present, 

2H++C0 3 --±^H 2 C0 3 , 

and to a small extent Ca ++ and Cl~ ions unite to form (easily 
ionized) CaCl 2 . The nearly complete removal of C0 3 ions 
allows the further ionization of CaC0 3 , and this change permits 
the passage of more solid CaC0 3 into solution. The quantitative 
relations are such that these changes continue until all CaC0 3 
has dissolved. Incidentally the H 2 C0 3 , which is unstable, dis- 
sociates, to a large extent, into water and C0 2 , 

H 2 C0 3 ^H 2 0+C0 2 f, 

and, as the latter is not very soluble, much of it passes off as 
a gas. 

The several reactions are shown in the following diagram: 

C0 2 (gas) 
It 
H 2 C0 3 ^H 2 0+C0 2 

It 

2HC1^2C1- + 2H+ 

CaC0 3 ^ CaC0 3 ^ Ca+++ C0 3 ~ 
(Solid) It 

CaCl 2 



Applications of the Ionic Hypothesis 293 

462. The Action of NaOH on NH 4 C1. — Another example of 
gas evolution, which, however, does not introduce any new 
principle, is found in the reaction 

NaOH+NH 4 Cl = NaCl+NH 4 OH, 

which takes place more or less completely when solutions of the 
two initial substances are mixed. This reaction was studied 
under Class II (426), where it was pointed out that it takes place 
nearly completely because NH 4 OH is a weak base. This base 
is also unstable, readily dissociating, thus, 

NH 4 OH^H 2 0+NH 3 t, 

and, since NH 3 is a gas, it will in part escape from the solution. 
The more concentrated the solution and the higher the tempera- 
ture the more completely will the NH 3 be evolved as gas. The 
loss of NH 3 from the solution promotes the dissociation of 
NH 4 OH, and this in turn favors a further shift in equilibrium 
from left to right in the main reaction. The various reactions 
and their relations are fully shown in the following diagram: 



NH 4 Cb 


NH 4 OH± 

It 
^C1-+NH 4 + 


NH 3 

It 
5 H 2 0+NH 3 


(gas) 


NaOH: 


^Na++OH- 

it 
NaCl 







463. The Factors Governing Gas Evolution. — The various 
factors which are favorable to gas evolution are very similar 
to those which were found to favor precipitation, although there 
are some differences aside from the fact that in the one case we 
are dealing with a gas and in the other with a solid product. If 
one of the products of a reaction is gaseous it will be given off 
from the solution, the more completely, the larger the propor- 
tion of it formed in the reaction and the less soluble it is. In 
these respects gas evolution is completely analogous to precipita- 
tion. Since all gases are less soluble at high than at low tempera- 
tures, gas evolution is always more complete the higher the 



294 Introduction to General Chemistry 

temperature. Gas evolution and precipitation differ in one 
very important additional respect: at a given temperature the 
M.S. of a precipitate has a fixed value, while that of a gas depends 
upon the pressure of the gas above its solution. In most cases 
the total solubility and therefore also M.S. of a gas is directly 
proportional to its (partial) pressure (Henry's Law, 126). If 
during gas evolution the partial pressure of the gas given off 
is kept down by removing the gas (as by blowing it away with a 
stream of air) as fast as it is liberated, its M.S. will be reduced to 
a vanishingly small value. In consequence the dissolved gas 
will be practically or even completely removed from the solution. 
Thus, in the reaction between NaOH and NH 4 C1, if a current of 
air is blown through the solution every trace of NH 3 will finally 
be removed, so that the reaction will be absolutely complete. 
The remaining solution will contain only common salt. The 
same result is attained if the solution is boiled, in which case the 
evolved steam takes the place of the air current. The high 
temperature also hastens the removal of the NH 3 . All reactions 
giving gases which follow Henry's Law may be driven to com- 
pletion by the continuous removal of the gas by means of a 
current of an inert gas or by steam produced when the solution 
is boiled. We have seen (455, 461) that the reason why a little 
soluble salt dissolves is the efficient removal from solution of one 
or both of its ions to form some new substance, which of course 
must be soluble or volatile if the resulting mixture is to be a 
solution. In the reactions studied so far the removal of ions 
has been accomplished by the formation of little ionized or 
little soluble substances. There are other ways of removing 
ions. These we shall take up later. We shall find that the 
dissolving of little soluble substances in question depend upon 
the same fundamental principle, and that these new cases differ 
from those studied in this chapter only in secondary ways, the 
means by which the ions in question are removed from the solu- 
tion (532, 560, 626). 

464. The Value of the Ionic Hypothesis. — In chapters xvii 
and xviii, we have applied the ionic hypothesis to the interpreta- 
tion of reactions between acids, bases, and salts and have seen 



Applications of the Ionic Hypothesis 295 

that this hypothesis leads to fairly simple explanations of a 
great variety of facts. Furthermore we have seen that if we 
know the degrees of ionization and the solubilities of the sub- 
stances concerned in any reaction we are able to predict what 
the result of the reaction will be. Herein lies the enormous 
practical value of the ionic hypothesis. 

In chapter xvii (412) we called attention to some of the 
glaring inconsistencies of the hypothesis; but we have also 
pointed out that the practical value of any hypothesis is not its 
truth but its usefulness. Having now, we hope, shown its use- 
fulness, we shall in later chapters consider whether it is true 
(chaps, xx, xxvii). 



CHAPTER XX 
ELECTROCHEMISTRY 

465. Introduction. — The present chapter will deal first with 
some of the marvelous developments of our knowledge of elec- 
tricity and matter during the last two decades. We now have 
good reason for believing that electricity like matter is of a 
granular or " atomic" nature. The grains or "atoms" of free 
electricity are all exactly alike and of the variety known as 
negative electricity. These grains are called electrons. Posi- 
tive electricity is not known in a free state; that is, it is only 
known as a positive charge on ordinary chemical atoms or 
larger masses of matter. 

466. The Granular Nature of Electricity; Electrons. — In 
chapter xvii (403) it was shown that Faraday's Law of Electro- 
chemical Equivalents leads directly to the conclusion that all 
univalent ions, in solution, carry equal charges of electricity (404) . 
The charge on a single univalent ion may be called a unit charge. 
Each bivalent ion has two unit charges, each trivalent ion three 
unit charges, etc. As early as 1874 Storey pointed out that these 
facts indicate that electricity is granular in nature, that each 
univalent atom is associated with one such granule to form a 
univalent ion, that a bivalent ion is an atom with two granules 
of electricity, etc. Furthermore a little later he proposed to 
call the quantity of electricity of a single granule an electron. 
This quantity is exceedingly minute. The common unit of 
quantity, one coulomb, is equal to more than a billion billion 
electrons. According to present-day usage the term electron 
means a single electronegative granule of electricity. 

467. Proof of the Existence of Electrons. — Although evidence 
was gradually accumulating during the last quarter of the nine- 
teenth century, it was not until more recently that positive 
evidence was obtained that electricity is granular and is made 
up of electrons. The crowning work was that of Professor 

296 



Electrochemistry 



297 



Robert Millikan, an American physicist who showed that when 
a very small sphere is charged with more and more electricity 
the quantity of electricity increases by small, equal additions, and 
not continuously. This is exactly what we should expect if the 
charge is made up of a small number of electrons. 

The spheres used were oil drops of microscopic size, not 
much larger, in fact, than the particles of dust that can be seen 
floating in the air when a beam of sunlight penetrates a dark- 
ened room. A drop was made visible by a beam of bright light 
and was viewed through a short-focus telescope. In still, dust- 
free air the drop fell, under the action of gravity, at a constant 
velocity that could be accurately measured. It is interesting 



t=z 




Fig. 75 



to note that although every body, however small, will fall with 
steadily increasing velocity in a vacuum, a very small body 
falls with constant velocity in air, owing to the viscosity of the 
latter. The drops used fell with a velocity of about one milli- 
meter per second. The principal parts of Millikan's apparatus 
are shown in Fig. 75 : M and N are parallel metal plates insulated 
from one another and connected through a switch to the terminals 
of a high-potential battery, B. The upper plate, M, can be 
charged positively and the lower one, N, negatively. A minute 
oil drop, D, is caused to fall into the space between M and N 
through a pinhole in the center of M, and its rate of fall is 
measured while M and N are uncharged, observations being 
made with a telescope, T. A minute negative charge is then 
given to the drop (in a way that need not be considered at 
present), and the plates M and N are charged. The drop is 



298 Introduction to General Chemistry 

now attracted by M and repelled by N, so that it moves upward. 
When it is close to M the electric field is switched off (S, switch; 
E, earth), and the drop is again allowed to fall, and its speed 
(time of fall) is again measured. With uncharged plates (field 
off) the drop falls at exactly the same rate, whether it is charged 
or uncharged. Next its speed upward is measured with the 
field on. This speed is always the same as long as the charges 
on M and N remain constant (constant field), and the charge 
on the drop is unchanged. But increase of negative charge on 
the drop increases the upward speed, and decrease of negative 
charge decreases the upward speed, the field remaining constant. 
The speed upward is a measure of the force of electrical attrac- 
tion by M and repulsion by N of the charged drop and is there- 
fore a measure of the charge on the drop. A drop could be made 
to make hundreds of trips up and down. The downward 
velocity (field off) was always the same; the upward velocity 
(field on) varied with the charge. The charge for each upward 
speed was found by a simple calculation. 1 It turned out that 
the charge on the drop was in every case a multiple by a whole 
number of the smallest possible charge on the drop. Thousands 
of observations were made in these experiments, but not an 
exception was found to the foregoing statement. This proves 
conclusively that electricity is granular in nature. It has been 
shown in other ways that the granules of electricity composing 
the charge on an oil drop are of the same magnitude as the unit 
charges of ions of electrolytes in solution. We may therefore call 
them electrons and say that all electric charges are made up of 
one or more electrons. In Millikan's experiments the oil drops 
used were observed to carry all possible charges from a single 
electron to over a hundred electrons; in no single case was a 
fraction of an electron found. The electron is therefore the 
smallest indivisible particle of electricity. 

468. The Nature of an Electric Current. — The relation of 
an electric charge to an electric current was first clearly estab- 
lished in 1876 by the American physicist Rowland, who showed 

*A popular account of Professor Millikan's work is given in his book, The 
Electron. Chicago: The University of Chicago Press, 191 7. 



Electrochemistry 299 

that when an electrically charged gilt disk was very rapidly 
rotated it produced the same sort of deflection on a magnetic 
needle as that due to a current of electricity flowing through a 
wire having the same position with respect to the needle as 
that occupied by the disk, Fig. 76. This experiment proved 
that a current of electricity is nothing but an electric charge 
in motion, just as a current of water is nothing but water in 
motion. 

469. The Electron Theory of Electric Currents. — If we 
accept the view that an electric current is an electric charge in 
motion, and also take into account the fact that an electric 
charge is merely an assemblage of electrons, we are at once led 
to the supposition that a current in a wire 
is only a stream of electrons passing through 
the wire. 

470. The Structure of an Atom. — If 
we think of a metal wire as made up of 
"solid," impenetrable atoms, it is not 
very reasonable to imagine that par- 
ticles of electricity (electrons) could pass 
through it. However, physicists and FlG 6 
chemists have in recent times come to 

the conclusion that an atom is by no means a homogeneous, 
solid lump, but that it is a rather complex structure, con- 
sisting largely or even wholly of negative electrons rotating 
in more or less circular orbits about a positively charged 
nucleus. The sum of the negative charges of the electrons 
is exactly equal to the positive charge of the nucleus, so that, 
as a whole, an atom has no excess of either kind of elec- 
tricity. The structure of an atom may be likened to that of 
the solar system, in which the sun corresponds to the nucleus 
and the planets to the surrounding electrons. The distances 
between the electrons composing an atom are probably large 
compared with the size of an electron, so that a stray electron 
might pass through an atom with the same ease that a comet 
passes through the solar system, or a bullet may pass through 
a flock of birds without striking any one of them. 




300 Introduction to General Chemistry 

471. How a Wire " Carries a Current." — If we think of a 
wire as made up of atoms of the sort here pictured, it is easy to 
see how a stream of electrons might pass through it. In a wire 
(not connected with any electrical source) some electrons are 
continuously becoming detached from their original atoms; 
these probably move through and among the atoms, occasionally 
replacing, for the time being, those that have been lost by other 
atoms. A metal always contains more or less of these free, 
wandering electrons, as well as an equal number of atoms which 
are deficient in electrons. When the terminals of a battery are 
joined by a wire, the positive pole of the battery attracts and 
the negative pole repels the free electrons of the wire. This 
causes a drift of electrons along the wire, and this drift of electrons 
constitutes the current in the wire. The progress of electrons in 
the direction of the drift is slow, a matter of a few centimeters 
per minute. The well-known fact that the effect of closing an 
electric circuit is felt almost instantaneously at a great distance 
(as illustrated, for example, by our everyday telephone experi- 
ence) is explained by the assumption that all the mobile electrons 
in the wires of the circuit move forward at the instant the circuit 
is closed. The case is just like that of drawing water from a 
supply pipe; water flows out the instant the faucet is opened, 
being pushed forward along the whole pipe by the water forced 
into the pipe by the pump. 

472. The Direction of an Electric Current. — Before the 
nature of an electric current had been discovered it was cus- 
tomary to consider that the current in the wire flowed from the 
positive to the negative pole. Since the drift or flow of electrons 
is in the opposite direction, there is danger of misunderstanding 
in speaking of the direction of the current. It is perhaps best 
to speak of the direction of the negative current, which is then 
the direction of drift of the electrons. 

473. Nonconductors of Electricity. — All metals are good con- 
ductors, but the non-metals are practically nonconductors or in- 
sulators. To account for this difference we have only to suppose 
that a non-metal, like sulfur, contains but very few free or mobile 
electrons and therefore has very little ability to carry a current. 



Electrochemistry 301 

474. Production of Electric Charges by Friction. — If a glass 
rod is rubbed with a piece of silk, the former takes on a positive, 
the latter a negative, charge. This is explained by assuming 
that a few stray electrons of the glass have been " wiped off" by 
the silk. The rubbing of the glass by the silk is of importance 
only in insuring intimate contact between the two. Another 
example of similar nature is found in the familiar electrification 
of the hair when combed with a hard-rubber comb in dry weather. 
Here the comb acquires a negative charge and the hair a posi- 
tive one. In general, when any two different substances are brought 
together they become electrified with opposite charges. This may be 
taken to mean that electrons accumulate in excess more easily 
on some kinds of matter than on others. 

475. Cathode Rays. — ^_____ \ 

When a high- voltage electric r^J^ ^'- ' : - ■'-■'- r ~-- r ''- - -~&* " " '""'"""' ' "f^ I 
current is passed through a ^-~-! N «-^--- - - '- ^- = -^: =x r. ".".'; ""_:;_ zM j 

Crookes tube, Fig. 77, which' K*i __ '/ 

is an evacuated glass bulb W 

having a metallic cathode, FlG 

C, and an anode, A, rays, 

known as cathode rays, are given off by the cathode and cause 
a yellowish-green fluorescence of the opposite end of the tube. 
These rays are readily stopped by a sheet of metal, as shown by 
the fact that a screen (in the form of a Maltese cross) casts its 
shadow on the glass. Even transparent substances like glass 
do not transmit the cathode rays any better than do metal sheets 
of comparable thickness. Extremely thin sheets of material like 
aluminum or gold leaf permit partial transmission of the cathode 
rays. 

476. X-Rays. — Cathode rays produce X-rays, also known as 
Roentgen rays, which radiate from any object struck by the 
former. A modern X-ray tube is shown in Fig. 78. This is a 
modified Crookes tube, intended for the use of large currents 
and the production of powerful X-rays. The cathode rays come 
from the specially constructed cathode, C, and strike a target, T, 
made of metallic tungsten, which metal is chosen because of its 
very high melting-point (3000 ). When the cathode rays are 



302 Introduction to General Chemistry 

stopped by the target, part of their energy is transformed into 
X-rays, and the balance appears as heat, so that the target 
becomes red, or even white, hot. Recent work has proved that 
the X-rays, which are very different from the cathode rays, are, 
like visible light, vibrations of the so-called luminous ether and 
differ from visible light in having wave-lengths about one- 
thousandth as great as the latter. 

477. The Nature of Cathode Rays. — The extensive investiga- 
tions of Sir William Crookes on cathode rays, during the seventies 
of the last century, led this famous English physicist to con- 
clude that these rays were matter in a highly rarefied or 




Fig. 78 

ultra-gaseous state, which he called a fourth state of matter 
(the other three states being solid, liquid, and gaseous). And 
in the light of our present knowledge of the real nature of these 
remarkable rays we must admit that Crookes's conclusion was 
substantially correct, although it was by no means the last word 
on the subject. 

It has long been known that cathode rays travel in a straight 
line in a vacuum, but that they may be deflected in an arc of a 
circle by a transverse magnetic field. The apparatus shown in 
Fig. 79 serves for lecture demonstration of this interesting 
phenomenon. A narrow beam of rays coming from the cathode 
and passing through a slit in a mica plate strikes along a screen 
covered with a specially prepared form of zinc sulfide, which 
becomes luminous in the line where it is struck by the rays. If 
now a horseshoe magnet is presented so that the N pole is above 



Electrochemistry 303 

the plane of the paper and the 5 pole below it, the beam is 
deflected to the position of the curved line. 

It is a well-known fact that a wire, free to move, is deflected 
by a magnetic field when a current is passed through it. The 
direction of deflection of the wire is determined by the direction 
of the current in the wire. The deflection of the cathode rays 
by a magnetic field indicates that the rays are electricity in 
motion, the direction of deflection corresponding to that of a 
stream of negative electricity coming from the cathode, which is, 
of course, the negative electrode. If we grant that the current 
in the wire leading to the cathode is, in reality, only a stream of 
negative electrons in the wire, we have only to suppose that these 



electrons do not stop on reaching the cathode, but shoot out from the 
surface of the latter and thus constitute the cathode rays. 

478. Proof that the Cathode Rays Are a Stream of Electrons. 
— The conclusive proof that the cathode rays are a stream of 
negative electricity (presumably electrons, since all negative 
charges consist of electrons) was furnished by the work of 
Perrin, a French physicist. Perrin's apparatus is shown in 
Fig. 80. It was a special form of Crookes tube having the 
cathode at C, the anode at A, and at B an insulated metal 
receiver, into which the cathode rays could be deflected by means 
of a magnet. This receiver was connected by a wire to an elec- 
troscope, capable of detecting any electric charge given to the 
box and determining its sign, whether positive or negative. 
When the cathode rays were started no charge passed into the 
receiver until the rays were magnetically deflected so as to fall 
into the receiver; then the latter acquired a large negative 



304 Introduction to General Chemistry 

charge. To guard against stray electric charges the receiver 
was surrounded by a metal shield connected to the earth, E. 
The experiments above described, together with many other 
facts, have led to the conclusion that the cathode rays are com- 
posed of negative electrons shot out from the cathode with high 
velocity. 

479. The Mass of an Electron. — An electron behaves as 
thought it had mass. In the first place we know that moving 
electrons have energy, since the cathode rays can produce light, 
heat, and X-rays, all of which are forms of energy. Since the 
kinetic energy of a moving body is proportional to the product 



Fig. 80 

of its mass and the square of its velocity, we can account for 
the energy of the cathode rays by assuming the electrons to 
have mass. Furthermore the fact that it requires an appreciable 
magnetic force to deflect the cathode rays and that the extent 
of the deflection (for rays of a given velocity) is proportional 
to the strength of the magnetic force is also evidence that elec- 
trons have mass. One of Newton's laws is to the effect that a 
moving mass continues in a straight line unless acted upon by a 
transverse force. Conversely, if a force is required to deflect 
a moving electron, we are warranted in assuming that the latter 
has mass. By methods that we cannot explain here it has been 
shown that the mass of an electron is about one eighteen-hundredth 
that of an atom of hydrogen. 

480. The Beta Rays of Radium. — The spectacular properties 
of radium have been brought to the attention of nearly everyone, 



electrochemistry 305 

whether he is a student of chemistry or not. Radium gives out 
three kinds of rays, the alpha, beta, and gamma rays. Of these 
the beta rays very closely resemble the cathode rays. Like 
cathode rays they are deflected by a magnetic field in a direc- 
tion which indicates that they too are a stream of electrons 
shot out with high velocity from the radium. Radium is, by 
all ordinary tests, an element. It resembles barium as closely 
as potassium resembles sodium. Here then is an element that 
spontaneously gives off negative electricity in the form of elec- 
trons shot out with great velocity. 

The alpha rays have been proved to be atoms of the element 
helium, He (atomic weight = 4), each of which carries a double 
positive charge. These rays are also shot out with high velocity. 
The gamma rays are identical with X-rays. 

481. The Disintegration Hypothesis. — The extraordinary 
behavior of radium has been satisfactorily explained by the 
disintegration hypothesis of Rutherford and Soddy. These 
scientists assumed that a radium atom is not a homogeneous 
solid particle but a very complex structure made up of electrons 
revolving rapidly in more or less circular orbits about a nucleus 
of positive electricity in the manner already described. It is 
further assumed that an atom of radium may become unstable 
and throw of a single electron {beta ray) or a larger mass {an atom 
of helium, which is an alpha ray), leaving behind an atomic residue 
of smaller mass and therefore smaller atomic weight. This hy- 
pothesis is in complete accord with all known facts concerning 
radium and radioactive phenomena. 

482. The Electrical Nature of Matter. — The study of radio- 
active substances, of which, in addition to radium, about thirty 
are known, has led to the conclusion that the atoms of all ele- 
ments, whether radioactive or not, are constructed on the same 
same plan as that of radium. According to this hypothesis the 
atom of one element differs from that of another only in the number 
and arrangement of the' electrons composing it. The mass of an 
atom is, at least in part, accounted for by the mass of the elec- 
trons composing it. All matter is considered to be of electrical 
origin. 



306 Introduction to General Chemistry 

483. The Nature of an Ion. — A single sodium ion is an atom 
of sodium having a single positive charge of electricity equal in 
quantity but opposite in sign to that of an electron. The 
simplest explanation of the difference between an ion and an 
atom of sodium is found in the assumption that the ion is an 
atom which has lost one electron. The atom was originally elec- 
trically neutral, because the positive charge of its nucleus was 
just equal to the sum of the negative charges of its surrounding 
electrons. If one electron is lost, the atom will have an excess 
positive charge just equal in magnitude to that of one electron. 
Since metallic atoms all form positive ions we conclude that all 
such atoms are able to lose electrons. Moreover, an atom of 
a univalent metal can lose but one electron and its ion will have 
a single unit charge, thus, 

Na(atom)->Na + + one electron. 

A bivalent atom can lose two electrons, 

Ca(atom)->Ca ++ + two electrons. 

A trivalent atom, such as that of aluminum, can lose three elec- 
trons, etc. 

Later work has shown that the ions are undoubtedly hydrated 
to some extent. The actual formula of sodium ion might be 
represented thus: 

Na(H 2 0) x +. 

The subscript x represents a small integer, probably 2 or 3. In 
practice we do not include the water in the formula, since in 
the first place the exact data necessary are wanting, and in the 
second place the relationships in our reactions seem to be satis- 
factorily represented without it. 

484. Valence. — The idea just presented leads to a simple 
explanation of valence (147, 183). The metals which form 
only positive ions do so by the loss of one or more electrons from 
each atom. The valence of an atom of a metallic element is 
determined by the number of electrons it has lost. 



Electrochemistry 



307 



A negative ion, such as Cl~ , is an atom which has taken up 
an extra electron. Atoms of metals do not take up extra elec- 
trons. Only the atoms of non-metallic elements behave in this 
way. The valence of a negative ion consisting of one atom corre- 
sponds to the number of electrons the atom has acquired. 

485. Theory of the Union of Sodium and Chlorine. — It is 
well known that sodium and chlorine unite very energetically 
to form NaCl. The simplest explanation of the cause of union 
is found in the assumption that an atom of sodium has a great 
tendency to lose an electron, and that an atom of chlorine has a 
great tendency to take up an extra electron. The violent reac- 
tion that we observe when we bring these two elements together 
is only the outward manifestation of the passage of electrons 
from the atoms of sodium into the atoms of chlorine. The 
residue of the sodium atom now has an excess of positive elec- 
tricity, while the chlorine atom with its extra electron is charged 
negatively. Since unlike electric charges attract each other, 
we may well assume that the two parts of the NaCl molecule 
are held together by electrical attraction. 

486. The Cause of Ionization. — If two insulated bodies are 
oppositely charged, Fig. 81, the force with which they attract 




Fig. 



each other is proportional to the product of their charges and 
inversely proportional to the square of the distance between 
them. There is, however, one additional factor that determines 
the strength of the attraction, and that is the nature of the 
surrounding medium. Usually this is air. If the medium were 



308 Introduction to General Chemistry 

glass instead of air the attraction would be only about one-third 
as great, other things remaining the same; but if the medium 
were water the attraction would be only one-eightieth as great 
as for air. If then we dissolve NaCl in water, the molecules are 
surrounded by a medium which lessens enormously the attract- 
ive force which holds their parts together; as the result, mole- 
cules will tend to fall apart, thus, 

NaCl->Na+ + Cl-. 

The positive part is the sodium ion, the negative part the chlorine 
ion. According to this explanation the molecule of salt before 
it ionizes is made up of two oppositely charged parts. These 
are not ordinary atoms, since the one has lost an electron which 
the other has gained. We ought to say that the NaCl mole- 
cule is made up of a sodium ion (Na + ) electrically bound to an 
ion of chlorine (Cl~). The act of ionization, which takes place 
when the salt is dissolved, is only the falling apart of the ions already 
present, on account of the great decrease in attractive force caused 
by the surrounding water. In other words, molecules are com- 
posed of bound ions, while in solution part of the ions are free. 
The ionization of all acids, bases, and salts is explained in pre- 
cisely analogous fashion to that in the case of NaCl. 

487. The Electronic Description of Electrolysis. — According 
to the electronic description of electrolysis, when an ion reaches 
an electrode it either gains electrons or loses them. Thus the 
positive ions Cu ++ , Pb ++ , H + , etc., each gain enough electrons 
to make them electrically neutral; while Cl~, I~, S~~, and 
0~~ each lose electrons and become free elements. 

488. The Displacement of Non-metals by One Another. — It 
will be recalled that chlorine acts on a solution of hydrobromic 
acid or any bromide, setting free bromine, thus : 

CL+ 2HBr->Br 2 + 2HCI. (259) 

Similarly bromine acts on iodides, as, for example: 
Br 2 +2KI->I 2 + 2 KBr. 



Electrochemistry 309 

Iodine acts on H 2 S, in solution, setting free sulfur: 

I 2 +H 2 S->S+2HI. . (339) 

The order in which the four above-mentioned elements displace 
one another is therefore as follows: 

CI, Br, I, S. 

Each will set free from its compounds any one following it. We 
may also include fluorine and oxygen in the series, and, since 
fluorine will displace any of the other elements mentioned, it will 
head the series. The position of oxygen 1 is determined by the 
fact that a H 2 S solution reacts with atmospheric oxygen to form 
free sulfur and water, 

2 +2H 2 S->2S+ 2 H 2 0, (339) 

and that a solution of HI also reacts with oxygen of the air 
(slowly) to form water and free iodine, 

2 + 4 HI^ 2l 2 + 2 H 2 0. (265) 

On the other hand, HBr solution is scarcely affected by oxygen 
gas, and HC1 solution not at all. Oxygen will therefore precede 
iodine and sulfur and follow bromine in the list. The whole 
displacement series is then as follows: 

F, CI, Br, O, I, S. 

489. Electronic Interpretation of Displacement. — If the re- 
action 

Cl 2 +2HBr±5Br 2 +2HCl 

takes place in very dilute solution, the two acids are nearly com- 
pletely ionized, and we may leave the H + ion out of considera- 
tion. The reaction in its simplest aspect is as follows: 

Cl 2 +2Br-^Br 2 + 2 Cl-. 

This means that each Br - ion loses an electron, which, passing 
into a chlorine atom, changes the latter into a Cl~ ion. Wo 
conclude that chlorine atoms take up electrons more readily than 



310 Introduction to General Chemistry 

do atoms of bromine. Considering next the six elements of the 
displacement series, we may say that fluorine has the greatest 
tendency to take up electrons, and sulfur the least, and that the 
tendencies of the other elements come in the order indicated in 
the list as given. Summarizing, we may say that of two elements 
of the above-mentioned, series the one whose atoms have the greater 
tendency to take up electrons will set the other free from its compounds 
with positive ions. 

490. The Displacement of Metals by One Another. — Strips 
of metallic zinc placed in solutions of salts of lead, copper, 
mercury, and silver will react as indicated by the following 

equations : 

Zn+Pb(N0 3 ) 2 ^Pb+Zn(N0 3 ) 2 , 
Zn+ CuS0 4 ->Cu+ZnS0 4 , 
Zn+HgCl 2 ^Hg+ZnCl 2 , 
Zn+ 2 AgN0 3 ->2Ag+Zn(N0 3 ) 2 . 

In other words, zinc displaces each of the above-mentioned metals 
from its salts. 

If strips of metallic lead are placed in solutions of salts of 
zinc, copper, mercury, and silver, no reaction takes place with 
the zinc salt; but the other three metals are set free, while the 
lead atoms pass into solution as positive ions. In similar 
fashion metallic copper sets free mercury and silver from their 
salt solutions, but it does not affect solutions of zinc or lead salts. 
Mercury displaces silver from its salts but has no action on salts 
of zinc, lead, or copper. Metallic silver will not displace from 
their salts any of the other metals just considered. The order 
of displacement of the five metals is therefore as follows : 

Zn, Pb, Cu, Hg, Ag. 

491. Electronic Interpretation of Metallic Displacement.— 

The action of zinc on solutions of copper salts may be represented 
in simplified form thus: 

Zn+Cu++->Cu+Zn++. 

This means that an atom of zinc gives up two electrons to an 
atom of copper. Since zinc displaces copper equally well from 



Electrochemistry 311 

solutions of all its simple soluble salts, we conclude that an 
atom of zinc has a greater tendency to lose electrons than has 
an atom of copper, but, since metallic copper displaces silver 
from any of its salts, thus, 

Cu+2Ag+-> 2 Ag+Cu++, 

we conclude that an atom of copper has a greater tendency to 
lose electrons than has an atom of silver. 

The order of the metals in the displacement series 

Zn, Pb, Cu, Hg, Ag 

is therefore the order in which they fall according to the decreasing 
ease with which their atoms tend to lose electrons. In the case of 
any two metals of the preceding list the one whose atoms have 
the greater tendency to lose electrons will set the other free from 
its compounds with negative ions. 

492. A More Complete Displacement Series of Metals. — 
Most of the familiar metals may be included in a single displace- 
ment series, which shows at the same time the tendencies 
of the atoms of the metals to lose electrons and so change 
into positive ions. The list is given in Table XIX. In this 

TABLE XIX 

Displacement Series of the Metals 

Potassium Nickel 

Sodium Tin 

Barium Lead 

Calcium Hydrogen 

Magnesium Copper 

Aluminum Mercury 

Zinc Silver 

Iron Platinum 

Cobalt Gold 

list (in which the second column follows the first) each metal 
tends to displace, or set free from its combination with negative ions, 
any element which follows it. 

Hydrogen has been placed in the list between lead and copper. 
Any metal above hydrogen in the series will react with a normal 



312 Introduction to General Chemistry 

solution of hydrochloric acid to set free hydrogen (at atmospheric 
pressure) and pass into solution as chloride. The metals follow- 
ing hydrogen in the list do not react readily, if at all, with hydro- 
chloric acid. The first four elements of the series react with 
water to set free hydrogen. Therefore metallic potassium 
placed in a solution of XaCl does not set free metallic sodium 
but causes the evolution of hydrogen. The order shown for 
the first four elements of the table is in fact that of their tend- 
encies to lose electrons as determined by means other than 
direct displacement. 

493. The Production of an Electric Current. — In the reac- 
tion between zinc and copper sulfate the essential change, as we 
have seen, is that represented by the equation 

Zn+ Cu ++ ->Zn ++ + Cu . 

We have said that this change is the result of the passage of 
two electrons from each atom of zinc into an atom of copper. 
Xow if this is true we ought to be able to get an available elec- 
tric current from the reaction; but if a piece of 
zinc is dipped into a solution of a copper salt, 
Fig. 82, no evidence of the production of such 
a current is to be observed. How indeed could 
we expect to detect the production of an electric 
current under the conditions pictured in Fig. &2? 
CuSOt If a passage of electrons occurred, it would be 
between the zinc rod and the layer of copper 
FlG g , ions in the solution surrounding the rod. and we 

could not readily detect such a current, much 
less make any use of it. If we wish to make this supposed 
current available for detection and use, we must so arrange the 
reacting substances that the Cu ~ ions are not directly in con- 
tact with the zinc rod. and then provide a wire for the transfer 
of electrons from the zinc rod to the copper ions. This can be 
done by arranging the four substances of the reaction 

Zn+ CuS0 4 ->ZnS0 4 -f Cu 






Electrochemistry 



3*3 



in the manner shown in Fig. 83 . Here we have a zinc rod dipping 
into a solution of ZnS0 4 in one beaker, and a copper rod dipping 
into a solution of CuS0 4 in the other beaker. A glass tube 
rilled with ZnS0 4 solution and loosely stoppered with cotton 
plugs forms a so-called salt bridge between the two beakers. If 




Fig. 83 

now the two rods are connected by wires to a galvanometer, a 
current is found to flow in a direction indicating the passage of 
electrons from the zinc rod through the wire (and galvanometer) 
to the copper rod. At the same time metallic copper begins to 
deposit on the copper plate, and metallic zinc begins to pass into 
solution. In fact, the reaction 

Zn+ CuS0 4 ^ZnS0 4 + Cu 

begins to take place just as soon as the metallic circuit is closed 
between the upper ends of the zinc and copper rods. No action 
occurs before the circuit is closed, and all action stops when the 
circuit is broken. 

494. The Mechanism of Current Production. — In detail the 
actions that occur with closed circuit are as follows : zinc atoms 
pass from the rod into the solution, each atom of zinc leaving 
behind two electrons and changing thereby into a Zn ++ ion. 
The electrons thus liberated flow through the wire to the copper 
rod in the CuS0 4 solution. Copper ions in contact with the 
copper rod take up two electrons each, being thereby changed 
into ordinary copper atoms. These latter adhere to the copper 
rod as a metallic coating. Fresh Cu ++ ions move up to the 
copper rod by diffusion, so that, as the ions in contact with the 
rod take on electrons and change into copper atoms, others move 
up by reason of their kinetic motion to take their places. On 
the other hand Zn ++ ions, newly formed at the zinc rod, diffuse 



314 Introduction to General Chemistry 

out into the solution. These changes tend to cause a deficiency 
of S0 4 ~~ ions about the zinc rod, and an excess of the same ions 
about the copper rod. The attraction between the excess of 
S0 4 ~~ ions, on the one hand, and the excess of Zn ++ ions, on 
the other, causes a migration of these ions in opposite directions 
through the solution and the salt bridge, and thus serves to 
maintain in each cubic centimeter of the whole solution as many 
negative as positive ions and thereby to keep the solution, as a 
whole, electrically neutral. 

495. The Function of the Salt Bridge. — The necessity of 
some sort of connection between the solutions of ZnS0 4 and 
CuS0 4 in the two beakers, Fig. 83, is obvious. If we remove 
the salt bridge, which in this case is a ZnS0 4 solution, the circuit 
is broken, and all action comes to a stop. By the use of the 
bridge we are able, by placing the CuS0 4 in a separate beaker, 
to keep it from coming in contact with the Zn rod. The use of . 
a metal wire in place of the salt bridge would apparently be a 
simpler plan, but it would not serve, because new products 
would be set free by electrolysis at each end of the wire. We 
could, however, use in the bridge, instead of the ZnS0 4 , a solu- 
tion of CuS0 4 or, in fact, of NaCl or almost any other salt. In 
case the bridge contains NaCl, the Na + ions serve in place of 
Zn ++ to carry the positive charge from the ZnS0 4 solution 
to the CuS0 4 solution, and the Cl~ ions to carry the negative 
charge in the opposite direction. 

496. Galvanic Cells. Electric Batteries. — A galvanic cell, 
or, as it is more popularly known, an electric battery, is any 
kind of apparatus by means of which an electric current is 
produced by chemical reactions. Dry batteries and storage bat- 
teries are, at present, the most familiar types. The first prac- 
tical form of the zinc-copper cell just studied was known as 
the Daniell cell; a later modification is known as the gravity 
battery. Properly speaking, the term battery means a group 
of cells, but this term is frequently used at present to mean a 
single cell. 

497. The Gravity Battery. — A gravity cell is shown in 
Fig. 84. It consists of thin sheets of copper surrounded by a 




Electrochemistry 315 

solution of copper sulfate in the lower half of the glass jar and a 

heavy zinc " crowfoot" surrounded by a zinc sulfate solution in 

the upper half. Attached to the copper sheets is an insulated 

copper wire. A new cell is set up by rilling the jar with water, 

placing the copper and zinc in position,, and adding more than 

enough solid CuS0 4 (blue vitriol or bluestone) to saturate the 

lower layer. No ZnS0 4 need be added; instead, 

twenty or thirty g. of NaCl are sprinkled into 

the water. The solution is not stirred. The 

CuS0 4 gradually dissolves, giving a saturated 

solution which soon fills the lower part of the 

cell. If now the insulated wire leading from 

the copper is connected to the zinc, a current 

flows through the wire, and the changes already 

described take place. The NaCl is used to _ 

r m Fig. 84 

make the water conduct the current prior 

to the formation of sufficient ZnS0 4 for this purpose. Until 

recently the gravity battery was used to operate all telegraph 

lines. 

498. Other Kinds of Galvanic Cells. — It is possible to make 
a cell that will give a current by the use of any pair of metals 
(not acted upon by water) , each surrounded by a solution of one 
of its salts. In each cell the experimental arrangement may be 
that shown in Fig. 83. 

499. Electromotive Force and Voltage. — A body at rest can 
be set in motion only by the action of a force (Newton's law). 
In a similar manner we assume that the current (stream of 
electrons) produced by a battery is the result of an electrical 
force called the electromotive force, E.M.F. The unit of 
E.M.F. is the volt (named after the pioneer electrical experi- 
menter Volta). The gravity cell has an E.M.F. of 1 . 1 volts. 

The farther the two metals forming the electrodes of a cell 
of any kind are removed from each other in the displacement 
series (492) the greater the E.M.F. of the cell. The reason for 
this is found in the fact that the metals heading the list give 
off electrons most readily (with greatest force). The order in 
the list represents, in fact, the relative force with which the 



316 Introduction to General Chemistry 

element loses electrons. The difference of such forces for the 
two metals of a cell is, for practical purposes, the chief deter- 
mining factor of the E.M.F. (voltage) of the cell. This difference 
of forces between the electrodes is also often called the potential 
difference of the electrodes. 

There is another important factor to be considered besides 
the nature of the reactions at the electrodes, and that is the con- 
centration of the ions in solution. For example, the more con- 
centrated the copper ions at the copper electrode the faster is the 
reaction carried on by these ions at a given temperature. Now 
the difference of potential at the terminals of a cell is a measure 
of the rate of the reaction in progress ; hence it will be increased 
or decreased by concentration changes in the solutions. To 
make careful comparison of the electromotive forces of cells the 
concentrations of the ions must therefore be taken strictly into 
account. However, in the series we are considering, no moderate 
variation of the concentrations of the ions from those found in 
the ordinary laboratory reagents (o.oi to 6N approximately) 
will produce results different from those described here, in the 
cases under consideration. The effect of the concentration of 
ions on cell potentials should be considered in an exact study of 
the latter subject. 

500. Electrical Energy. — Electrical energy always depends 
on two factors, voltage and quantity of electricity. The unit 
of electrical energy is the joule, named after J. P. Joule, the 
celebrated English scientist, whose work on the mechanical 
equivalent of heat was discussed earlier (370). One joule is the 
amount of energy produced when a quantity of one coulomb of elec- 
tricity flows through a conductor, the ends of which have a potential 
difference of one volt. In general, joules = voltsX coulombs. 
For example, if a gravity cell of 1 . 1 volts E.M.F. delivers 10 
coulombs, the electrical energy produced is 1.1X10=11 joules. 
Since the joule is an energy unit, its value is expressible in other 
energy units.. Careful experiment has shown that 

1 joule= 10,200 g.cm., 
1 joule = . 24 calorie, 
1 calorie = 4. 1 8 joules. 



Electrochemistry 



3i7 



It is electrical energy which a consumer pays for and uses. 
The same number of electrons go back to the positive pole of a 
battery as leave the negative pole, but they lose energy in so 
going. The energy which the electrons give up may be liberated 
as heat or may be converted into work by means of devices like 
the motor. 

501. Electronic Explanation of Oxidation-Reduction Reac- 
tions. — The action of chlorine on ferrous chloride in solution 
(173, 332) is a simple, typical example of an oxidation-reduction 

reaction, 

2FeCl 2 +Cl 2 ->2FeCl 3 . 

This reaction in dilute solution may be represented by the simpli- 
fied equation 

2Fe++ + Cl 2 ->2Fe+++ + 2Cl-. 

The ferrous ion, Fe ++ , which is the reducing agent, is oxidized 
to Fe +++ by the chlorine atom, which is the oxidizing agent. 
This is explained by assuming that the Fe ++ ion (which is an 
iron atom that has already lost two electrons) gives up a third 
electron, which, passing into the CI atom, changes the latter 
into a Cl~ ion. Thus we see that the oxidation of the Fe ++ ion 




Fig. 85 

consists in its loss of an electron; and the reduction of the CI atom 
consists in its gain of an electron. 

502. Oxidation-Reduction Cells. — The transfer of electrons 
which occurs in the reaction just studied can be made to yield 
an available electric current quite as readily as that which 
takes place in the reaction between metallic zinc and copper 
sulfate (493). We may demonstrate this fact by means of the 
arrangement shown in Fig. 85. Platinum electrodes are placed 



318 Introduction to General Chemistry 

in each of two beakers, one of which contains the FeCl 2 solution, 
the other the Cl 2 solution (together with some FeCl 3 or NaCl to 
make the solution conduct). A salt bridge joins the two solu- 
tions. Wires from the electrodes are connected with a galvano- 
nometer, which shows the passage of a current in a direction 
indicating a flow of electrons in the wire from the electrode in the 
FeCl 2 solution to that in the Cl 2 solution. The platinum elec- 
trodes serve as carriers of electrons into and out of the solutions. 
Platinum is superior to any other metal except gold for this 
purpose, because of its very slight tendency to pass into solution 
as ions. 

503. Further Examples of Oxidation-Reduction Reactions. — 
Oxidation-reduction reactions are very common. They may 
all be interpreted in terms of electron transfers, as the following 
additional examples will illustrate. Ferric sulfate is reduced by 
zinc according to the equation 

Fe 2 (S0 4 ) 3 +Zn->2FeS0 4 +ZnS0 4 . ( 335 ) 

The simplified ionic equation is 

2Fe++++Zn^2Fe++-fZn++. 

Each atom of zinc loses two electrons and changes into a Zn ++ 
ion; these two electrons are taken up, one by each Fe +++ ion, 
which is thereby changed to a Fe ++ ion. The zinc, which loses 
electrons, is the reducing agent and is oxidized by ferric ions, 
which gain electrons and are thereby reduced to ferrous ions. 

The action of ferric salts and soluble iodides is illustrated by 
the following reaction: 

2FeCl 3 +2KI^2FeCl 2 -f-2KCl+I 2 , 

or in simplified form by 

2 Fe+++ + 2i— >2Fe++ + I 2 . 

A closely analogous reaction takes place in the reduction of 
ferric salts by hydrogen sulfide : 

2FeCl 3 +H 2 S^2FeCl 2 +2HCl+S. 



Electrochemistry 319 

The simplified equation is 

2 Fe+++ + S--->2Fe++ + S. 

The electronic explanation of each of the foregoing reactions can 
easily be made by the student. 

Other more intricate oxidation-reduction reactions, which 
will require a somewhat more extended discussion, will be taken 
up in subsequent chapters. 

In all oxidation-reduction reactions transfers of electrons 
occur; and in all cases the atom or ion which is oxidized loses 
one or more electrons, and the atom or ion which is reduced gains 
one or more electrons. If an ion does not change its charge or 
its composition in the course of a reaction it is neither oxidized nor 
reduced. 

504. The Oxidation and Reduction of Metals. — When a 
metal passes into solution its atoms take on positive charges. 
This means that each atom of a metal loses one or more elec- 
trons when it changes into an ion. Since we have defined 
oxidation as the loss of electrons (501) we can therefore say 
that when a metal changes into its ions it is oxidized. For 
example, in the reaction 

Fe+ CuS0 4 ->FeS0 4 + Cu, 
which we may write in simplified form thus, 

Fe+Cu++->Fe++ + Cu, 

we say that the metallic iron is oxidized to ferrous ions, and the 
copper ions are reduced to metallic copper. We have already 
seen that the further oxidation of Fe ++ to Fe +++ involves a 
loss of one additional electron. 

505. The Oxidation and Reduction of Non-metals. — When 
a non-metal (chlorine for example) passes into solution, its atoms 
take on electrons. We say, therefore, that in such a case the 
element is reduced. Conversely we say that its ions are oxidized 
when by loss of electrons they are changed to atoms of the 
element. 



320 



Introduction to General Chemistry 



506. Oxidation-Reduction Potentials. — Every oxidation- 
reduction reaction can by suitable arrangement be made to 
yield an electric current. The E.M.F. (voltage) of an oxidation- 
reduction cell is the measure of the force with which the reaction tends 
to take place. The stronger the oxidizing tendency of the oxidiz- 
ing agent and the stronger the reducing tendency of the reducing 
reagent the greater the E.M.F. of the cell. A systematic study 
of such cells has shown that all oxidizing and reducing agents may 
be arranged in a series in the order of their decreasing oxidizing 
tendencies and increasing reducing tendencies. 

507. Oxidation and Reduction by Means of the Electric 
Current. — We have shown that oxidation and reduction are 
capable of producing electric currents. There now remains 
to show that an electric current can accomplish oxidation and 
reduction. Two beakers, Fig. 86, are fitted with platinum 



FeCI, 




~ if 



Fig. 86 



electrodes and joined with a salt bridge, and in one is placed a 
solution of FeCl 3 , and in the other HCL Upon passing a current 
from a battery of two dry cells so connected that the electrode 
in the FeCl 3 will be the cathode, it will be found that the FeCl 3 
is reduced to FeCl 2 , while at the same time HC1 is oxidized to 
free chlorine at the anode. The explanation is as follows: The 
battery sends a steady stream of electrons through the wire to 
the cathode; one electron passes from the latter into each 
Fe +++ ion coming in contact with it, changing the Fe +++ 
into Fe ++ . At the anode Cl~ ions coming in contact with 
this electrode give up to it their electrons and change thereby 
into ordinary CI atoms. The latter then unite in pairs to form 
molecules, aggregates of which soon form bubbles that escape 
into the air. In the solution Fe +++ ions are attracted by and 
migrate toward the cathode,. while Cl~ ions are attracted by and 



Electrochemistry 321 

migrate toward the anode. Thus the transfer of electricity from 
one electrode to the other in the solution is accomplished by means 
of the moving ions, while in the wire we have a stream of electrons 
set in motion by the battery. A great variety of other oxidations 
and reductions in solution can be accomplished by means of the 
electric current. In fact, since we may consider the change of 
a metal into its ions as an oxidation of the former, and the reverse 
change a reduction of the ions, we may go farther and say that 
all processes of electrolysis result in oxidation and reduction. The 
anode is the seat of oxidation, since it takes up electrons; the 
cathode is the seat of reduction, because it furnishes electrons. 
These statements apply to all electrolyses irrespective of whether 
the substances formed or liberated at the electrodes are elements 
or compounds. 

508. The Conversion of Chemical Energy into Electrical 
Energy. — The production of heat by a chemical reaction has 
been explained (373) as being due to the conversion of chemical 
energy of the reacting substances into heat energy. If metallic 
zinc is placed in a solution of copper sulfate so that the reaction 

Zn+CuS0 4 ->ZnS0 4 -fCu 

takes place without the production of an available electric 
current, the quantity of heat liberated is 50,100 calories for one 
symbol weight of zinc. If the same amount of zinc reacts with 
copper sulfate in a gravity cell, 2X96,500 coulombs of elec- 
tricity are delivered into the circuit at an E.M.F. of 1.09 volts. 
The electrical energy produced is 2X96,500X1.09 = 210,400 
joules. Since 1 calorie = 4. 18 joules, 210,400 joules = 210,400 
-7-4.18 = 50,300 calories. Thus we see that electrical energy 
equivalent to 50,300 calories is produced in a cell, instead of 
50,100 calories of heat produced when the same amounts of the 
substances react directly, without the production of a current. 
The small excess of energy produced in a cell is accounted for by 
the fact, established by experiment, that this amount of energy 
is absorbed as heat from the surroundings as the cell operates. 
Similar results are observed in the energy production of all other 
galvanic cells. The electrical energy produced by any cell is 



322 Introduction to General Chemistry 

equal to the chemical energy liberated or lost, plus or minus an 
additional amount of energy — plus if heat is taken up from the 
surroundings and minus if it is given out to the surroundings. We 
may consider a galvanic cell or battery merely as a device for 
converting chemical energy into electrical energy. 

509. Conversion of Electrical Energy into Chemical Energy. — 
In all processes of electrolysis electrical energy is used up in the 
production of new chemical substances, and we may conclude 
at once that the electrical energy used is changed to chemical 
energy. 



CHAPTER XXI 
NITROGEN AND AMMONIA 

510. Introduction. — Many facts about nitrogen and two of 
its important compounds, ammonia and nitric acid, have already 
been discussed. We may recall that about four-fifths by volume 
of air is uncombined nitrogen, which is left in a nearly pure state 
as a colorless, odorless gas when air is freed from oxygen (15). 
Pure nitrogen is obtained by passing ammonia over red-hot 
copper oxide (84) : 

2NH 3 + 3 CuO->3Cu+ 3 H 2 0+N 2 . 

The symbol weight of nitrogen is 14, and since 22.4 liters of the 
gas at o° and 76 cm. weigh 28 g., the formula is N 2 (75). This 
means that a molecule of nitrogen consists of two atoms (215). 

Ammonia, NH 3 , is a colorless gas which can be made by the 
union of nitrogen and hydrogen, by the action of electric sparks, 

N 2 + 3 H 2 -> 2 NH 3 . (298) 

It has a very penetrating odor and dissolves easily in water, 
forming at the same time some ammonium hydroxide: 

NH 3 +H 2 O>NH 4 0H. (91) 

Ammonia and ammonium hydroxide may be completely elimi- 
nated from a solution by boiling the latter. 

Ammonium hydroxide is a base which neutralizes acids to 
form salts, among which we have studied the chloride NH 4 C1 
(92), the nitrate NH 4 N0 3 (105), the sulfate (NH 4 ) 2 S0 4 (101), 
and the acid sulfate NH 4 HS0 4 (101). 

Ammonium hydroxide is a very much weaker base than 
sodium hydroxide (409, 429); the latter substance reacts with 
ammonium salts to set free NH 3 , as illustrated by the following 
equation: 

NH 4 Cl+NaOH->NaCl+H 2 0+NH 3 . (426) 



324 Introduction to General Chemistry 

511. The Occurrence of Nitrogen. — Besides the occurrence 
of free nitrogen in air, this element is also found in nature as a 
constituent of many compounds, among which are the familiar 
substances ammonia and the nitrates of sodium and potassium. 
Nitrogen is also an essential constituent of the proteins, which 
constitute the bulk of all kinds of flesh. Proteins are also present 
in plants, particularly in their seeds. Cereals are especially 
rich in nitrogenous matter. In the course of the decay of animal 
and vegetable matter the nitrogen is changed first into XH 3 and 
finally into nitrates. 

512. The Element Nitrogen. — Xitrogen was discovered by 
Rutherford, professor of chemistry in Edinburgh, in 1772. 
Lavoisier, who was the first to recognize it as an element, called 
it azote. Its English name, nitrogen, suggested by Chaptal, 
indicates that the element can be made from saltpeter. KX0 3 , 
for which the Greek name was nitron. 

513. The Preparation of Nitrogen. — So-called atmospheric 
nitrogen, obtained by removing oxygen, carbon dioxide, and 
water from air, is not pure and cannot easily be freed from its 
residual impurities, amounting to about 1 per cent, which con- 
sist of chemically inactive gases, chiefly argon. Atmospheric 
nitrogen is prepared on a technical scale by the Linde process, 
in which air is first liquefied and the more volatile nitrogen 
distilled from the less volatile liquid oxygen. Pure nitrogen 
can be made not only from NH 3 and CuO but also in several 
other ways. One of the best of these is the following: 20 g. 
of XH 4 C1 and 25 g. of sodium nitrite, a salt having the formula 
XaX0 2 . and 50 c.c. of water are placed in a flask (Fig. 87) and 
heated gently, whereupon the following reaction occurs: 

NH 4 Cl+NaN0 2 ->NaCH-N 2 +2H 2 0. 

This reaction takes place in two stages, of which the flrst is 
XH 4 Cl+XaX0 2 ^XaCl+XH-X0 2 . 

The salt XH.X0 2 . ammonium nitrita, is unstable and decomposes 
at once into X 2 and H 2 0, thus: 

XH 4 X0 2 ->X 2 +2H,0. 



Nitrogen and Ammonia 



325 




514. The Properties of Nitrogen. — Nitrogen is a colorless, 
odorless gas, as might be inferred from the fact that four-fifths 
by bulk of the air is free nitrogen. The gas does not support 
combustion, nor does it burn in the ordinary sense of the term. 
However, nitrogen unites with oxygen, under certain conditions, 
to form the oxides NO and N0 2 . These very important reactions 
will be considered in detail in chapter xxii, as by their means 
nitric acid can be made from the nitrogen of the air. 

Nitrogen unites di- 
rectly with several ele- . Q 
ments, mostly metals, 
and especially at high 
temperatures, to form 
nitrides, among which 
calcium nitride, Ca 3 N 2 ; 
magnesium nitride, 
Mg 3 N 2 ; and aluminum Fig. 87 
nitride, A1N, may be 

mentioned. Nitrides usually react readily with water to form 
ammonia and the hydroxide of the metal: 

AlN+3H 2 0->Al(OH) 3 -f-NH 3 . 

A reaction of great technical importance takes place when 

nitrogen gas is passed over calcium carbide, CaC 2 (49), at a 

white heat: 

CaC 2 +N 2 ->CaCN 2 +C. 

The product, CaCN 2 , is calcium cyanamid; it is also called 
nitrolime. It is a valuable fertilizer and is also used in the manu- 
facture of ammonia and cyanides by methods which will be 
described later. 

515. The Assimilation of Nitrogen by Plants. — All plants 
require for their growth nitrogen in some form, in addition to 
other things. Fertile soils contain considerable nitrogen in the 
form of compounds, among which are ammonium salts, nitrates, 
and decaying vegetable and animal matter. Such compounds 
supply plants with their required nitrogen, since in general plants 
can assimilate only combined and not free nitrogen. A few 



326 Introduction to General Chemistry 

plants, however, among which are clover, alfalfa, and other 
legumes, have the power of utilizing uncombined nitrogen from 
air which has permeated the soil. This they do through the 
medium of root nodules, which are bacterial growths, whitish 
in color and often as large as peas. Free nitrogen is taken up 
by the bacteria in the nodules and changed into compounds, 
which then serve as a source of nitrogen for the plant. A large 
part of the nitrogen taken up by plants is concentrated in their 
seeds as complex compounds called proteins. Wheat, for 
example, contains 10 to 15 per cent of proteins. 

516. The Sources of Ammonia. — Although ammonia occurs 
in very minute amounts in the air, in natural waters, and in the 
soil, being formed by the decay of animal and vegetable matter, 
such occurrences are not a practical source of supply. In earlier 
times ammonia, then called spirits of hartshorn, was made by 
heating bones, hoofs, horns, and other animal matter with lime 
(calcium oxide). At present ammonia is produced in large 
amount as a by-product in the distillation of coal for the pro- 
duction of fuel and illuminating gas and coke. Coal usually 
contains 1 or 2 per cent of combined nitrogen, and when heated 
to a high temperature with exclusion of air, the nitrogen com- 
pounds present are decomposed, with the formation of ammonia, 
which pass'es off with the gas. The ammonia is absorbed in 
dilute sulfuric acid, giving a solution which upon evaporation 
yields crystals of ammonium sulfate; and this salt, when heated 
with lime, yields free ammonia: 

(NH 4 ) 2 S0 4 +CaO->CaS0 4 +H 2 0+2NH 3 . 

517. The Physical Properties of Ammonia.— For the sake 
of ready reference we may summarize here the more important 
data regarding ammonia. It is a colorless gas of penetrating 
odor; one liter at o° and 76 cm. weighs o. 772 g. ; and 22.4 liters 
weigh 17 grams. Upon being strongly compressed ammonia 
condenses to a colorless liquid. Liquid ammonia is an important 
article of commerce; it is marketed in steel cylinders of about 
two hundred pounds capacity. At 20 the liquid has a vapor 
pressure of 8.4 atmospheres. If this pressure is released the 



Nitrogen and Ammonia 327 

liquid boils; that is, it changes into gaseous NH 3 , and its tem- 
perature falls to —S3° (519)? which is its boiling-point under one 
atmosphere pressure. 

At 20 water dissolves about seven hundred times its volume 
of NH 3 gas, part of it going to form the hydroxide NH 4 OH, 
which is a weak or slightly ionized base. Pure ammonium 
hydroxide of commerce has a specific gravity of o . 90 and con- 
tains 28 per cent by weight of NH 3 . 

518. The Uses of Ammonia. — Ammonia is a substance of 
fundamental importance. It is made and consumed in immense 
quantities, and its uses are numerous. In the household it is 
added to water used in washing. In laboratories and chemical 
works it is indispensable. Ammonium salts, particularly the 
sulfate, are used in enormous quantities in fertilizers. The 
manufacture of nitric acid from ammonia has become since the 
beginning of the war a matter of vital importance; this subject 
will be taken up in the next chapter (570) . The use of ammonia 
in making ice may now be considered in some detail, as it is not 
only of commercial importance but also of much scientific 
interest. 

519. The Manufacture of Artificial Ice. — It will be recalled 
that the change of water into steam at ioo° takes place with 
the absorption of 540 calories of heat per gram (115); and 
further that " every pure liquid has a latent heat of evaporation." 
The latent heat of evaporation of liquid ammonia at its boiling- 
point, — 33°, is 258 calories. This means that liquid NH 3 
released from pressure passes into gas with an absorption of 258 
calories of heat per gram. This heat is taken from the ammonia 
and the vessel containing it, and if this vessel is immersed in 
water the latter is quickly frozen. Since the conversion of 1 g. 
of water at o° into ice requires an absorption of 79 calories, 
enough heat is absorbed by the evaporation of 1 g. of NH 3 to 
freeze about 3 g. of water. 

The practical application of these principles to the manu- 
facture of artificial ice is illustrated diagrammaticaUy by Fig. SS. 

The pump, A, draws ammonia gas from the pipes immersed 
in the tank, O, and compresses the gas sufficiently so that when 



328 



Introduction to General Chemistry 



the latter is cooled by running water in the condenser, B, 
liquid ammonia is formed, and collects in receiver, C. From 
C liquid ammonia flows into the pipes that run through the 
tank, 0. In the pipes the ammonia evaporates, thus cooling 
the brine (freezing-point — 20 ) with which is filled. The 
water to be frozen is contained in cans of 200 pounds or more 
capacity, immersed in the brine. 




Fig. 88 

520. The Synthesis of Ammonia. — We have already learned 
that NH 3 can be made from nitrogen and hydrogen by passing 
an electric spark through the mixed gases (298). The electric 
spark is effective mainly as a source of heat : 

3 H 2 +N 2 ->2NH 3 . 



This reaction reaches a state of equilibrium when but very little 
NH 3 has been formed, for the reason that the reverse action, 
the decomposition of NH 3 into H 2 and N 2 , takes place very 
easily. At a temperature of 8oo° equilibrium is reached when 
only one molecule of NH 3 remains undecomposed out of every 
10,000 taken. At first thought it would seem to be quite 
hopeless to try to make use of this reaction as a means of manu- 
facturing NH 3 from H 2 and N 2 , because of the very minute pro- 
portion of these gases that unite before the reaction reaches 
equilibrium. In order that a process based on this reaction 
shall be a practical success, substantially all of the H 2 and N 2 
taken must be converted into NH 3 . The big problem then is 
to obtain from the H 2 and N 2 taken the other 9,999 molecules 
of NH 3 . On account of the great technical importance of the 
subject and as a beautiful illustration of the way in which 



Nitrogen and Ammonia 329 

fundamental chemical principles (laws) are used in working 
toward desired results, we shall consider this problem in detail. 
521. The Effect of Temperature on the Equilibrium. — We 
have already learned (288) that for any reaction the proportions 
of the substances present when equilibrium is reached are 
different at different temperatures. With change of tempera- 
ture the equilibrium of any reaction is shifted in a way that can 
be definitely predicted from .its heat of reaction. If heat is 
given out when a substance is formed, heat will be absorbed 
when the same substance is decomposed. When we raise the 
temperature of a system in equilibrium, the state of equilibrium 
shifts in such a way that heat is absorbed (288, 367). The heat 
of formation (361) of NH 3 is 12,200 calories: 

3H 2 +N 2 ->2NH 3 H-2Xi2,2oo cal. 

Therefore the higher the temperature the smaller the propor- 
tion of NH 3 present in the equilibrium mixture, and vice versa, 
as shown by the following table, in which is given the percentage 
of NH 3 at various temperatures and one atmosphere pressure: 

Temperature Per Cent of NH 3 

8oo° 0.01 

6oo° 0.05 

400 0.48 

It is evident that the lower the temperature the more favorable 
the result; but a serious practical difficulty is encountered if 
one tries to make NH 3 at low temperatures; the union of H 2 
and N 2 proceeds the more slowly the lower the temperature, so 
that at 400 it would require days for equilibrium to be reached. 
In this respect this reaction is like all others; the speed of reac- 
tion, other things being equal, is slower the lower the tempera- 
ture. The experimenter then faces this dilemma: at a high 
temperature the reaction gives very little NH 3 ; at a low tem- 
perature it goes too slowly. Is there no way to surmount this 
difficulty? Is there no way to hasten a reaction except by 
increasing the temperature? Yes, there is, and this by means 
of a catalytic agent! 



33 o Introduction to General Chemistry 

522. Catalytic Agents for the Ammonia Reaction. — As illus- 
trations of cases of catalysis we may recall that CuCl 2 catalyzes 

the reaction 

4 HCl+0 2 ->2H 2 0+ 2 C1 2 ; (239) 

that Mn0 2 promotes the decomposition of KC10 3 (306); and 
that platinum black causes H 2 and 2 to unite rapidly at room 
temperature (303). Therefore it will not be surprising to learn 
that several catalytic agents were found which hasten the 
ammonia reaction enormously; among such were the metals 
iron, manganese, and uranium. By the use of these and prob- 
ably other substances, the nature of which has not yet been 
disclosed by the discoverers, the ammonia reaction is caused to 
reach equilibrium rapidly at 400 to 500 ; but even at 400 
only one-half of 1 per cent of ammonia is formed. In the next 
paragraph we shall see how the equilibrium can by another means 
be shifted still farther in the desired direction. 

TABLE XX 

Percentages of Ammonia at Various 

Temperatures axd Pressures 



t,„„ . Percentage at Percentage at 

Temperature £ Atmos £ here [ IOO Atmosphere; 



8oo° 
7 00° 
6oo° 
500 



O.OI I.I 

0.02 2.1 

O.05 . 4.5 

O.I3 IO.S 



523. The Effect of Pressure on the Equilibrium. — In chapter 
xiii (287) we learned that u the effect of increase of pressure on 
any system in equilibrium is in all cases to shift the equilibrium 
so as to favor the formation of substances occupying a smaller 
volume." In the reaction 

3 H 2 +N 2 ->2NH 3 

three volumes of hydrogen and one volume of nitrogen, if com- 
pletely united, will yield two volumes of ammonia (77), from 
which we conclude that an increase of pressure will shift the 
equilibrium so as to favor the formation of more XH 3 . Table 
XX gives the results. 



Nitrogen and Ammonia 331 

It has been found practicable to work at pressures of 100 
atmospheres (1,400 lb. per square inch) and even higher, and 
thus greatly to increase the proportion of NH 3 formed. It now 
remains to show how all of the H 2 and N 2 taken may be converted 
into NH 3 . 

524. The Effect of Removing One Product of a Reaction. — 
In the discussion of the reaction 

NaCl+H 2 S0 4 ->NaHS0 4 +HCl (289) 

it was shown that if the HC1 continuously escapes as a gas the 
reaction will go to completion from left to right. In the reaction 
between hydrogen and nitrogen equilibrium results from two 
opposing changes, each of which continues in the equilibrium 
mixture : these are the union of H 2 and N 2 on the one hand and 
the dissociation of NH 3 on the other. If the NH 3 formed can 
be removed from the mixture in some systematic way, the H 2 
and N 2 will finally be completely united. We shall now briefly 
describe the process as actually carried out. 

525. The Manufacture of Synthetic Ammonia. — The nitrogen 
used in making synthetic ammonia might be obtained from the 
air by the Linde process (513), and the hydrogen could be made 
by the electrolysis of water, if these methods were not too 
expensive. A practical method of obtaining the required mixture 
of N 2 and H 2 is said to have been worked out, in which air and 
steam are passed over heated coke. The carbon comprising the 
coke unites with the oxygen of both the air and the steam. The 
C0 2 and CO formed are then removed, leaving N 2 and H 2 . 

The apparatus used in the production of NH 3 is represented 
diagrammatically in Fig. 89. The mixture of N 2 and H 2 under 
high pressure, say 100 atmospheres, enters at A and passes into 
the chamber B, which holds the catalytic agent and is surrounded 
by an electric heater C. The reaction 

3 H 2 +N 2 ->2NH 3 

takes place to the extent of perhaps 8 to 10 per cent in B. and 
the mixture of the three gases passes on to the vessel D, sur- 
rounded by the refrigerating vessel E. The high pressure and 



332 



Introduction to General Chemistry 



low temperature cause nearly all of the NH 3 to condense in D 
to the liquid state, while the H 2 and N 2 , being much less readily 
liquefied, pass out and by means of the pump, F, are conveyed 
by a pipe back to the reaction vessel, B. Liquid NH 3 is removed 
from D through the valve at the bottom of D. 




c^q 



Fig. 89 



526. Ammonia from Cyanamid. — Cyanamid, CaCN 2 , is 
made by the action of nitrogen on calcium carbide at a red heat 
(514): * 

CaC 2 +N 2 ->CaCN 2 +C. 



When cyanamid is treated with steam under high pressure 
the following reaction takes place: 

CaCN 2 + 3 H 2 0->2NH 3 +CaC0 3 . 

Cyanamid has been manufactured at Niagara Falls, Canada, since 
1909; and since 191 5 considerable quantities of ammonia have 
been made at this place by the reaction described above. 
Recently the U.S. government has built an immense cyanide 
ammonia plant at Muscle Shoals, Alabama. It now appears 
possible that this process will be economically superior to 
that by which ammonia is made directly from nitrogen and 
hydrogen. 

527. The Chemistry of Ammonia. — One of the most familiar 
reactions of ammonia is that by which white fumes are formed 



Nitrogen and Ammonia 333 

when ammonia gas comes in contact with a volatilized acid. 
The fumes are, of course, fine particles of the resulting salt: 

NH 3 +HC1->NH 4 C1, 

NH 3 +HN0 3 ->NH 4 N0 3 . 

In water solution we have the same type of reaction with acids, 
resulting in the formation of ammonium salts. A consideration 
of the equilibrium equations 

NH 3 +H 2 O^NH 4 OH^NH 4 ++OH- 

will show that we may represent the reaction of aqueous ammonia 
solution with acids as the result of the union of the hydrogen 
ion of the acid with the hydroxyl ion of the base to form water, 
and the consequent displacement of the equilibrium to the right ; 
or we may represent the reaction as the union of NH 3 with H + 
of the acid to form NH 4 + , and the consequent displacement of 
the equilibrium to the left. The net result of either reaction is 
the same, and very likely the truth is that both routes are 
actually followed. The final solution contains water and the 
highly ionized ammonium salt. 

In these reactions of ammonia, either as a gas or in solution, 
the nitrogen atom suddenly and very easily takes on one posi- 
tive and one negative atom or radical. So striking is the ease 
with which these reactions occur that we think that the nitrogen 
atom was ready to take on the new groups and so must carry 
a positive and a negative charge, balanced against each other, in 
addition to the three negative charges by which it holds the three 
hydrogen atoms, thus, =±=NH 3 . From what we know of the 
complicated structure of an atom it seems plausible to assume 
the existence of a positive and a negative charge on the same 
atom. The reaction of aqueous ammonia with an acid may then 
be represented as follows: 

±NH 3 +H+ + C1-^NH 4 ++C1-^NH 4 C1. 

A compound which, like ammonia, contains an atom which is 
not exerting its full valence is said to be unsaturated. There 
are many other unsaturated compounds. 



334 Introduction to General Chemistry 

Though the above-mentioned reactions are the most char- 
acteristic ones shown by ammonia, we have seen one case of a 
different kind, namely the reducing action of ammonia on copper 
oxide (52, 84) : 

2NH 3 + 3 CuO->N 2 +3H 2 0+3Cu. 

An even more vigorous oxidation of ammonia occurs when a 
mixture of ammonia and air is passed over red-hot platinum. 
Under the catalytic influence of platinum, nitric acid is formed: 

NH 3 + 20 2 ^HN0 3 + H 2 0. 

This reaction is the basis of one of the great processes for manu- 
facturing nitric acid from the air (570). 

When ammonia gas is passed over heated sodium the follow- 
ing reaction takes place: 

2 Na+2NH 3 ->2NaNH 2 +H 2 . 

The products are hydrogen and sodium amide. The latter sub- 
stance is completely decomposed by water, thus : 

NaNH 2 +H 2 0->NaOH+NH 3 . 

Ammonia reacts with other metals, such as magnesium, forming 
nitrides (514), compounds which are also hydrolyzed by water 
to form ammonia and the hydroxide of the metal. 

528. Ammonium Salts. — Ammonium salts closely resemble 
the corresponding salts of potassium. They are easily soluble 
in water and crystallize well from solutions. None of the salts 
are colored excepting those with colored acid ions. 

529. Dissociation of Ammonium Chloride. — Ammonium 
chloride may be completely volatilized at temperatures much 
below red heat. When the molecular weight of the salt was 
calculated on the basis of determinations of the density of the 
vapor formed at about 300 , the value found was half a formula 
weight. Apparently a full formula weight of ammonium 
chloride, if still a gas, at o° and 76 cm. pressure would occupy 
not 22.4 but 44.8 liters. To some chemists sixty to eighty 
years ago it seemed impossible to co-ordinate these facts with 



Nitrogen and Ammonia 335 

Avogadro's hypothesis, and they accordingly abandoned the 
latter. 

A simple experiment gives the clue, however, to the explana- 
tion of the enigma. If a bit of dry NH 4 C1 is placed in a dry test 
tube and cautiously heated, a piece of moist red litmus paper 
held in the mouth of the tube turns blue, thus indicating the 
formation of free NH 3 . Elaborate experiments have shown 
that the vapor of NH 4 C1 dissociates at 300 almost completely 
intoNH 3 andHCl: 

NH 4 C1^NH 3 +HC1. 

Since the molecules of ammonia are much lighter than those of 
hydrogen chloride they travel (diffuse) more rapidly at the same 
temperature (197). Hence, although the two gases are released 
at the same instant, the ammonia molecules are the first to reach 
the litmus paper in the mouth of the test tube. Upon cooling, 
the constituents unite again to form solid NH 4 C1. Therefore 
NH 4 C1 vapor has half the density expected, for the reason that 
the number of molecules present is practically double, the 
original, because each NH 4 C1 molecule gives one molecule of 
NH 3 and one of HC1. Thus Avogadro's law covers the case 
perfectly. Many other apparent exceptions to Avogadro's law 
are known, but in every case a satisfactory explanation has been 
found in complete harmony with the law. 

530. Dissociation of Other Ammonium Salts.— Several other 
ammonium salts also dissociate into the acid and ammonia. 
Thus solid ammonium bicarbonate smells strongly of ammonia, 
at ordinary temperatures, by reason of the following dissociation : 

NH 4 HC0 3 ±5NH 3 +H 2 0+C0 2 . 

Ammonium sulfide, formed by the union of ammonia with 
hydrogen sulfide, dissociates at room temperatures to form 
ammonia and hydrogen sulfide: 

(NH 4 )HS^NH 3 +H 2 S. 

Ammonium hydroxide itself, which may be considered as the 
product of the union of ammonia and the very weak acid water, 



336 



Introduction to General Chemistry 



is stable in crystalline form only at a very low temperature (91). 
In general we find that the weaker the acid the more unstable 
is the ammonium salt. It is perhaps well to point out that dis- 
sociations of the kind just considered are entirely distinct from 
the splitting up of salts, acids, and bases into ions in solution, in 
that the new molecules formed by the dissociation of a gas or 
vapor are not electrically charged. Furthermore the dissociation 
products are of different composition in the two cases, as is 
illustrated by Table XXI. 

TABLE XXI 



Substance 


Electrolytic 
Dissociation 


Gaseous 
Dissociation 


NH4CI 

NH4HCO3.... 

NH 4 HS 

NH4OH 


NH 4 + ci- 
NH4+ HC03- 

NH 4 + HS- 
NH 4 + OH- 


NH 3 HCl 
NH 3 H 2 C0 2 
NH 3 H 2 S 
NH 3 H 2 



531. Other Compounds of Nitrogen and Hydrogen. — Nitro- 
gen and hydrogen form several other compounds besides NH 3 ; 
among such hydrazine, N 2 H 4 , and its hydrate, N 2 H 5 OH, a basic 
substance, may be mentioned. Hydroxylamine, NH 2 0H, is 
also a base. Hydrazoic acid, HN 3 , is a violently explosive 
liquid. It is an acid and forms salts like NaN 3 , sodium azide, 
and AgN 3 , silver azide; the latter is extremely explosive. The 
ammonium salt NH 4 N 3 and the hydrazine salt N 2 H S N 3 may be 
considered as having the formulae N 4 H 4 and N 5 H S respectively. 

532. The Solubility of Silver Chloride in Ammonia Solu- 
tion. — Silver chloride, AgCl, as we have seen (169), is an almost 
insoluble white salt which is easily obtained by adding any 
chloride solution to a silver salt solution; for example, 

NaCl+AgN0 3 ->AgCl+NaN0 3 . 

This reaction goes nearly to completion by reason of the very 
small solubility of silver chloride, so that precipitation con- 
tinues until the concentrations of both Ag + and CI"" ions are 
very small (452). 



Nitrogen and Ammonia 337 

If pure silver chloride is stirred with water, only very little 
dissolves; in fact, about 1.3 mg. per liter. We have, then, 

Solid Dissolved 

AgCl^AgCl^Ag++Cl-. 

By far the greater part of the dissolved substance is present as 
Ag + and Cl~ ions. 

If a solution of ammonia is added to the AgCl and water the 
solid dissolves very easily, giving a clear, colorless solution. 
This solution is a good electrical conductor, while, as we have 
already learned, ammonium hydroxide conducts very poorly. 
These facts may be demonstrated by the use of the apparatus 
described in sec. 384. If the cell is first rilled with ammonia 
solution, it is found that the current which passes is insufficient 
to light the lamp. As soon as AgCl is dissolved in the ammonia 
solution the lamp glows brilliantly. As a matter of fact the 
new solution is as good a conductor as solutions of equal con- 
centrations of most salts. 

. 533. The Silver Ammonium Ion, Ag(NH 3 ) 2 + . — It has long 
been known that dry AgCl unites with dry NH 3 gas to form 
solid compounds. It is possible that one or more of these is 
present in the solution formed when AgCl dissolves in a solution 
of ammonia; but if so, what are the ions, if any, of the new silver 
compounds? We have already learned (397) that the composi- 
tion of the ions of a solution is best investigated by means of 
migration experiments. If we electrolyze a solution of AgCl 
dissolved in ammonia, using a U-tube like that shown in Fig. 47 
(397)? we find that both Ag + and NH 3 migrate toward the 
cathode or negative electrode, while the chloride ion migrates in 
the opposite direction. Moreover, we find that the Ag + and 
NH 3 migrating toward the cathode are in the proportion of one 
atom of Ag to two molecules of NH 3 as represented in the 
formula Ag(NH 3 ) 2 , so that it would appear that this last formula 
represents the composition of the positive ion. We are therefore 
led to conclude that the following reaction first takes place, 

AgCl+ 2 NH 3 ^Ag(NH 3 ) 2 Cl, 



338 Introduction to General Chemistry 

and that the compound so formed ionizes very readily thus : 
Ag(NH) 2 Cl^Ag(NH 3 ) 2 ++Cl-. 

The silver ammonium ion, Ag(XH 3 ) 2 + , is called a complex ion. 

534. The Stability of the Ag(NH 3 ) 2 ~ Ion. — It has been proved, 
by an electrical method which need not be considered here, that 
in a solution of silver chloride in ammonia only an exceedingly 
small number of Ag^~ ions are present. This means that the 
complex ion, Ag(NH 3 ) 2 + , is very stable and is dissociated only 
very slightly into Ag + and 2XH3, thus: 

Ag(NH 3 )+^Ag++2NH 3 . 

535. Why Silver Chloride Dissolves in Ammonia. — We can 

now explain why AgCl dissolves in ammonia solution. Let us 
consider the following formulation of the reaction: 

Solid Dissolved 

AgCl±»AgCteAg++Cl- 

2NH3 



r- 



Ag(NH 3 ), 



The larger part of the silver chloride dissolved by pure water is 
present as Ag + and Cl~ ions. If XH 3 is added, the Ag + ions 
unite with it nearly completely to form Ag(XH 3 ) 2 ~. This 
reaction greatly reduces the concentration of Ag + and therefore 
causes more solid AgCl to pass into solution. If sufficient XH 3 
is present these changes go on until all the solid AgCl has dis- 
solved. In the solution we have the salt silver ammonium 
chloride. 

Consideration of the equation 

Ag++2NH 3 ^Ag(NH 3 ) 2 + 

shows that by adding an excess of XH 3 the concentration of Ag f 
will be still further decreased. If we wish to dissolve AgCl com- 
pletely it is necessary to add a little more than two formula 
weights of XH 3 for one of AgCl in order to depress the con- 
centration of Ag + sufficiently. The dissolving of silver chloride 



Nitrogen and Ammonia 339 

in ammonia is exactly similar to the dissolving of silver acetate 
in nitric acid, except that in the latter case the acetate ion was 
suppressed, while in the former the Ag + ion was suppressed. 

536. The Effect of Acids on Silver Ammonium Ion. — If a 
strong acid, nitric acid for example, is added to silver ammo- 
nium chloride solution, silver chloride appears as a precipitate. 
Evidently the nitric acid has caused an increase in the con- 
centration of Ag + ion, and union of the latter with the abund- 
ant chloride ion has followed. This reaction is explained by 
saying that NH 3 is converted into NH 4 N0 3 by the HN0 3 . 
This the acid does by virtue of the hydrogen ion which it 
furnishes. The latter unites with the free ammonia as follows: 

NH 3 +H+^NH 4 +. 

Some union of the latter ion with the nitrate ion to form NH 4 N0 3 
also takes place, but the critical change is that of ammonia to 
ammonium ion. The removal of NH 3 allows the dissociation of 
the silver ammonium ion to proceed to completion: 

Ag(NH 3 ) 2 +->Ag+ + 2NH 3 . 

537. Other Silver Ammonium Compounds. — If ammonium 
hydroxide is added to silver nitrate solution a little at a time, 
brown silver hydroxide is seen to precipitate but eventually to 
redissolve, just as the ammonia added exceeds the proportion 
of two molecules for each molecule of silver nitrate present : 

AgN0 3 +NH 4 OH->NH 4 N0 3 +AgOH; 
AgOH+ 2 NH 3 ->Ag(NH 3 ) 2 OH. 

But silver ammonium hydroxide is a strong base, and as fast as it 
forms it displaces ammonium hydroxide from its salt, NH 4 N0 3 , 

NH 4 N0 3 +Ag(NH 3 ) 2 OH->NH 4 OH+Ag(NH 3 ) 2 N0 3 , 

so that the final equation for the formation of silver ammonium 
nitrate from ammonia and silver nitrate is 

AgN0 3 + 2NH 3 ->Ag(NH 3 ).N0 3 . 



340 Introduction to General Chemistry 

In the same way a solution of silver ammonium sulfate can 
be made by the addition of aqueous ammonia to silver sulfate. 
All the silver ammonium salts are found to be highly ionized, 
just as are potassium and ammonium salts. 

538. Other Complex Ions.— Ammonia unites with many 
other metal ions to form complex ions; for instance, with cupric 
ion it forms the brilliant blue copper ammonium ion, and 
with nickel ion it forms the less highly colored blue nickel 

ammonium ion. 

Cu+++ 4 NH 3 ->Cu(NH 3 ) 4 ++ 

Ni+-++ 4 NH 3 ->Ni(NH 3 ) 4 ++. 

These complex ions form salts such as Cu(NH 3 ) 4 S0 4 and 
Ni(NH 3 ) 4 (N0 3 ) 2 . The chemistry of these ions parallels that of 
the silver ammonium ion. The ammonium ions are the first 
which we have called complex ions. But as a mattter of fact 
we have been dealing with a great number, for example, nitrate, 
carbonate, phosphate, sulfate, and hydroxyl ions. In the double 
decomposition reactions studied, these ions have shown no sign 
of any dissociation into smaller parts corresponding to the 
dissociation of the silver ammonium ion into silver ion and 
ammonia. However, there is evidence that these secondary 
dissociations do exist, but to a very much smaller degree than 
in the case of a complex silver ion. In other words, these ions 
are very much more stable than the latter. We shall find later 
other examples of moderately stable complex ions like those of 
ammonia. 



CHAPTER XXII 
NITRIC ACID AND THE OXIDES OF NITROGEN 

539. Nitric Acid, HN0 3 . — That nitric acid is a substance of 
great importance will be apparent at once when it is known that 
it is an indispensable agent in the manufacture of explosives and 
dyesturls and in addition is used extensively in a great many 
other ways. We have already learned (104) that nitric acid can 
be made from sodium nitrate, Chile saltpeter, and sulfuric acid: 

NaN0 3 +H 2 $0 4 ->NaHS0 4 +HN0 3 . 

If a larger proportion of NaN0 3 is used, the following reaction 
can also occur if the temperature is high enough: 

2 NaN0 3 +H 2 S0 4 ->Na 2 S0 4 + 2 HN0 3 . 

Until recently the only commercial method of making nitric 
acid was by means of these two reactions. In practice the 
action takes place in large cast-iron stills, which are but slightly 
attacked by the two acids as long as water is. not present. The 
nitric acid distils and is usually condensed in vessels made of 
fused quartz, which is not acted upon by HN0 3 , or in Duriron (an 
alloy of iron and silicon) , which is almost unaffected by the acid. 

540. Chile Saltpeter, NaN0 3 . — It is a remarkable fact that 
there is but one known source of sodium nitrate of sufficient 
magnitude to be of practical importance: this consists of enor- 
mous deposits found in a desert region in the mountains of Chile. 
These deposits form a stratum averaging five feet in thickness 
and covering about six hundred square miles. The value of the 
saltpeter exported amounts to three-fourths of the total exports 
of Chile, and the export duty on it is the chief source of revenue 
of the country. 

Potassium nitrate is found in limited quantities in India, but 
the deposits are far too small to be of importance for the manu- 
facture of nitric acid. 

341 



34 2 Introduction to General Chemistry 

541. Properties of Nitric Acid. — Pure anhydrous (water- 
free) nitric acid is a colorless liquid of density 1.52; it boils at 
7 8°, undergoing some decomposition into water and oxides of 
nitrogen. The pure acid of commerce contains only 68 per cent 
of HNO3, the balance being water. This acid has a density of 
1.42. When dilute nitric acid is boiled the residual portion 
grows more concentrated, until it reaches a density of 1.42 and 
68 per cent purity. Acid of this concentration then distils at 
the constant temperature of 120 . Acid more concentrated 
than 68 per cent boils at lower temperatures, and as the boil- 
ing continues the temperature rises, while the residual acid 
becomes less concentrated, until finally, at 120 , 68 per cent 
acid is left. 

Nitric acid unites with water to form two crystalline hydrates, 
HN0 3 -H 2 and HN0 3 - 3 H 2 0, which melt at -38 and -18 
respectively. 

Nitric acid is one of the most highly ionized acids ; it therefore 
has all the properties characteristic of a strong acid. Its salts 
are all soluble in water — most of them very soluble. Nitrates 
of strong bases (Na, K, Ca, Ba, Mg, and Ag) give neutral solu- 
tions, while solutions of nitrates of weak bases are acid in reaction, 
owing to hydrolysis (436) . Other important properties demand 
treatment in separate paragraphs. 

542. Nitric Acid as an Oxidizing Agent. — In addition to its 
action as an acid, nitric acid also acts as a powerful oxidizing 
agent. It will be recalled that oxidation and reduction always 
go hand in hand (327) ; the oxidizing agent is reduced, and the 
reducing agent is oxidized. Most reducing' agents react readily 
with nitric acid, being themselves oxidized as the result of the 
reaction. The reduction of HN0 3 gives the compounds N0 2 , 
NO, N 2 0, N 2 , or NH 3 , according to the reducing agent used. 
The subject is best approached after one has become familiar 
with the properties of the oxides. We shall therefore first take 
up the study of the latter and then return to the discussion of 
the action of nitric acid as an oxidizing agent (557). 

543. Nitric Oxide, NO. — Nitric oxide is a colorless gas 
almost insoluble in water; it is most easily made by the 



Nitric Acid and Oxides of Nitrogen 343 

action of dilute HN0 3 on copper. We might expect the 

reaction to be 

Cu+2HN0 3 ->Cu(N0 3 ) 2 +H 2 , 

but copper has virtually no tendency to displace hydrogen from 
solutions of acids. We find that the reaction actually takes 
place according to the equation 

3Cu+8HN0 3 ^3Cu(N0 3 ) 2 +4H 2 0+2NO. 

The apparatus shown in Fig. 87 (513) is used. About 20 g. of 
copper in the form of wire, clippings, turnings, etc., are placed 
in a 250 c.c. flask, and 60 c.c. of concentrated nitric acid mixed 
with 60 c.c. of water are added. The wash bottle contains dilute 
NaOH. 

The reaction in the flask starts at once, and the vessel is soon 
filled with a brown-colored gas. After a few minutes this dis- 
appears, and in its place is a faintly colored gas, nitric oxide, 
which is impure. When purified by passing through the wash 
bottle it proves to be colorless. It may now be collected over 
water in the cylinder set in the pneumatic trough. 

544. Direct Union of Nitrogen and Oxygen. — The student's 
general chemical experience and observations will have already 
led him to the conclusion that nitrogen does not readily combine 
directly with oxygen; for if the reaction 

N 2 +0 2 ->2NO 

occurred very easily it would be impossible for the mixture of 
N 2 and 2 composing air to exist without union taking place. 
Nevertheless the passage of electric sparks through air leads to 
the formation of a small amount of nitric oxide; but the reaction 
soon reaches equilibrium because of the decomposition of nitric 
oxide into its constituents under the action of electric sparks. 
The passage of a flash of lightning through air produces a little 
nitric oxide. 

Although the union of nitrogen and oxygen under the influ- 
ence of electric sparks was discovered by Cavendish in 1766, it 
had only a scientific interest for chemists up to very recent times. 



344 Introduction to General Chemistry 

The recent application of this reaction to the manufacture of 
nitric acid will be taken up later (566). 

545. The Properties of Nitric Oxide, NO.— Nitric oxide is a 
colorless gas nearly insoluble in water and slightly heavier than 
air. One liter weighs 1 . 34 g. and 22.4 liters 30 g., agreeing with 
that calculated from the formula NO. Nitric oxide is very 
stable, not being decomposed into its elements except at high 
temperatures. 

When nitric oxide comes in contact with air it turns brown. 
The same change takes place when nitric oxide is mixed with half 
its volume of pure oxygen. The reaction takes place thus : 

2NO+0 2 -> 2 N0 2 . 

The brown gas, N0 2 , is called nitrogen tetroxide. This is the 
brown gas that first appears when nitric oxide is being made, 
since the latter unites with the 2 present in the flask to form 
brown N0 2 . It is interesting to note that a little heat is given- 
out when NO and 2 unite. 

546. Combustion in Nitric Oxide. — The proportion of oxygen 
in nitric oxide is about two and one-half times as great as in air, 
and we might expect that this gas would support combustion very 
readily; but this is not the case. Burning sulfur or a burning 
candle is extinguished if brought into a jar of nitric oxide. A 
mixture of hydrogen and nitric oxide does not explode with an 
electric spark. On the other hand, a few substances will burn 
in nitric oxide if strongly heated; thus an iron wire heated to 
incandescence by an electric current takes fire in this gas and 
burns, forming an oxide and setting nitrogen free. Briskly 
burning phosphorus continues to burn when plunged into a jar 
of nitric oxide, forming phosphorous pentoxide and nitrogen, 

4 P+ioNO->2P 2 5 +5N 2 . 

The vapor of carbon bisulfide, CS 2 , an easily combustible 
and very volatile liquid, gives with nitric oxide a mixture that 
readily burns with a bright light. By way of caution it may be 
pointed out that carbon bisulfide is a very dangerous liquid, 



Nitric Acid and Oxides of Nitrogen 345 

because of the readiness with which it catches fire. It must be kept 
away from all flames and heated objects. 

547. Other Ways of Making Nitric Oxide.— Nitric oxide is 
formed by the action of dilute nitric acid on many other metals 
besides copper. It is also formed when various oxidizable sub- 
stances react with dilute HN0 3 as illustrated by the following 
equation : 

2HN0 3 +6FeS0 4 +3H 2 S0 4 ^3Fe 2 (S0 4 ) 3 +2NO+4H 2 0. 

By means of this reaction very pure nitric oxide can be made. 

548. Nitric Oxide and Ferrous Sulfate. — When nitric oxide 
is passed into a solution of FeS0 4 a dark-colored solution is 
formed containing the compound FeS0 4 *NO. When this solu- 
tion is heated the compound is decomposed, and very pure 
nitric oxide is given off. The formation of a nearly black solu- 
tion with FeS0 4 serves as a good test for nitric oxide, and in con- 
sequence also for nitric acid, by reason of the fact that nitric 
acid is reduced to nitric oxide by FeS0 4 in the presence of H 2 S0 4 , 
the ferrous sulfate being oxidized to ferric sulfate, Fe 2 (S0 4 ) 3 , at 
the same time. 

549. A Test for Nitric Acid and Nitrates.— If 2 or 3 c.c. of a 
mixture of nitric acid or any nitrate and concentrated sulfuric 
acid are placed in the bottom of a test tube, and a solution of 
ferrous sulfate is cautiously added in such a way that the two 
solutions do not mix, a dark-brown 'ring will be formed at the 
junction of the two layers. The brown layer contains the com- 
pound FeS0 4 -NO. 

550. Nitrogen Tetroxide, N0 2 . — We have learned that the 
brown gas nitrogen tetroxide is formed by the direct union of 
nitric oxide and oxygen, thus: 

2N0+0 2 ^2N0 2 . 

It is also formed in several other ways, such as by the action of 
concentrated HN0 3 on many metals. For example, with con- 
centrated HN0 3 and copper we have 

Cu+4HN0 3 ->Cu(N0 3 ),+ 2NO a + 2H 2 0. 



346 Introduction to General Chemistry 

Dilute HNO3 and copper give NO instead of N0 2 (543). 
The cause of this difference is found in the behavior of the two 
oxides of nitrogen toward H 2 and HN0 3 , as represented in the 
equation 

3 N0 2 +H 2 0^ 2 HN0 3 +NO. 

With much warm water this reaction takes place nearly com- 
pletely from left to right; but the reaction is reversible, since 
nitric oxide acts almost completely on an excess of concentrated 
nitric acid to form nitrogen tetroxide and water. Therefore, in 
the action of nitric acid on copper, nitric oxide will be given off 
if the acid is very dilute, and nitrogen tetroxide if it is very con- 
centrated; with an acid of intermediate concentration a mixture 
of the two oxides will result. 

Nitrogen tetroxide is also prepared by heating the nitrates 
of heavy metals such as copper, lead, mercury, silver, etc. (565). 

551. The Physical Properties of Nitrogen Tetroxide.— 
Nitrogen tetroxide is a reddish-brown gas having a peculiar, 
disagreeable odor. It is dangerously poisonous and may pro- 
duce fatal results after a day or two when inhaled in quantities 
which seem inconsequential at the time. It is easily condensed 
to liquid form by cooling it with ice; at — io° liquid nitrogen 
tetroxide solidifies to a mass of nearly colorless crystals. The 
liquid formed by the melting of these crystals is also nearly 
colorless at temperatures near the melting-point ; as the tempera- 
ture is raised the liquid becomes first yellow, then orange, and 
finally boils at 22 , giving a gas or vapor of light reddish-brown 
color. At higher temperatures the gas becomes much darker in 
color. 

552. The Density of Nitrogen Tetroxide.— The formula N0 2 
corresponds to a formula or molecular weight of 46 and a gas 
density 1 . 59 times that of air. The actual density which at 
room temperature is very much greater than this decreases 
with rise of temperature and reaches the value 1 . 59 only at 
140 . The exact results are shown in Table XXII. 

The very remarkable changes in density and color with 
change of temperature are thought to be due to the formation 



Nitric Acid and Oxides of Nitrogen 



347 



at low temperatures of double molecules having the formula 
N 2 4 , the equation being 

2 N0 2 ^N 2 4 - 

If we accept this view and also suppose N 2 4 to be colorless and 
N0 2 dark reddish-brown we get a very satisfactory explanation of 
all the known facts. The form N 2 4 exists almost pure in the 
crystals at — io°; in the liquid state some dissociation to N0 2 
has taken place, giving a yellow or at higher temperature an 
orange-colored liquid. 

TABLE XXII 



Temperature 


Density 
Air = i 


Calculated 
Molecular Weight 


Per Cent of N0 2 


Per Cent of N2O4 


2 5 ° 


2.50 
I.78 

1 59 


72 
52 
46 


43 

87 

IOO 


57 

13 




97 


140 







In Table XXII the percentages of N0 2 and N 2 4 at various 
temperatures are given, as calculated from the densities. At 
140 the dissociation into single molecules of N0 2 is practically 
complete. In accord with this explanation it is actually found 
that the density (corrected of course for thermal expansion) 
remains constant above 140 . 

In chemical equations we shall continue to write the formula 
N0 2 for nitrogen tetroxide, but it will be understood that much 
of the gas, at room temperature, is in the form of N 2 Q 4 . 

553. Nitrites and Nitrous Acid. — A very interesting reaction 
takes place when N0 2 is passed into a solution of NaOH. The 
gas is completely absorbed, and the solution upon evaporation 
yields crystals of two salts — sodium nitrate, NaN0 3 , and sodium 
nitrite, NaN0 2 . The reaction occurs according to the equation 

2 N0 2 + 2NaOH->NaN0 3 +NaN0 2 +H 2 0. 

Sodium nitrite is a salt of technical importance, particularly 
in connection with the manufacture of certain dyestuffs and 
medicinal chemicals. It is the salt of an acid, HXO... called 



348 Introduction to General Chemistry 

nitrous acid. The free acid itself is very unstable and its dilute 
solution easily decomposes, thus: 

3 HN0 2 ->HN0 3 + 2NO+H 2 0. 

On the other hand the nitrites of sodium and potassium are 
stable in water solution. They are also resistant to heat. As 
a matter of fact they are usually made by the decomposition 
of the corresponding nitrates under heat (565). Ammonium 
nitrite, however, decomposes in hot-water solution, forming 
nitrogen and water (513). 

All nitrites are easily oxidized to nitrates. Solutions of these 
reagents undergo this change simply on exposure to air and 
consequently are never free from nitrates. 

Nitrous acid is an oxidizer strong enough to change hydriodic 
acid into iodine and water, but not strong enough to act similarly 
with hydrobromic acid: 

2HI+ 2 HN0 2 ->I 2 + 2NO+ 2H 2 0. 

Since nitrous acid is easily reduced to nitric oxide it will give 
the ring test (549), as the latter depends on the formation of 
this substance. 

554. Nitrogen Trioxide, or Nitrous Acid Anhydride. — An 

oxide having the composition N 2 3 is found to form white crystals 

at — 103 . At room temperatures it can exist only in minimal 

amounts in equilibrium with its decomposition products NO 

andN0 2 : 

N 2 3 ^NO+N0 2 . 

The decomposition of nitrous acid yields products which appar- 
ently come from the dissociation of this gas, so that the latter 
may be called the anhydride of nitrous acid (313) : 

2 HN0 2 ^H 2 0+N 2 3 - 

555. Nitrogen Pentoxide, the Anhydride of Nitric Acid. — 

Nitrogen pentoxide, N 2 5 , is a rather unstable, colorless crystal- 
line substance which is formed from pure anhydrous (water-free) 
HN0 3 by the action of P 2 O s , thus: 

2HN0 3 +P 2 5 ->2HP0 3 +N 2 5 . 



Nitric Acid and Oxides of Nitrogen 349 

The N 2 s formed is separated from the involatile metaphosphoric 

acid, HPO3, by careful distillation and condensation in a vessel 

cooled with ice. Nitrogen pentoxide decomposes readily into 

N0 2 and 2 , thus: 

2 N 2 5 ->4N0 2 +0 2 . 

It also unites vigorously with H 2 to form HN0 3 , 
N 2 O s +H 2 0^2HN0 3 . 

556. Nitrous Oxide, N 2 0, Laughing Gas. — Of the five oxides 
of nitrogen, nitrous oxide, N 2 0, is the only one met with outside 
of chemical works and laboratories. It is the well-known sub- 
stance laughing gas, used by all dentists as a mild anaesthetic. 

This gas is made practically by heating ammonium nitrate, 
which decomposes easily according to the following equation: 

NH 4 N0 3 ->N 2 0+ 2 H 2 0. 

If the substance is heated too strongly the reaction takes place 
explosively. In fact, NH 4 N0 3 is being used extensively as an 
explosive. Therefore great care must be taken to heat the substance 
cautiously. 

Nitrous oxide is a colorless gas, having a faint, not unpleasant 
odor. It is somewhat soluble in water but is usually collected 
over warm water. The gas is condensed to liquid form by 
sufficient pressure. Liquid N 2 boils at — 90 . This liquid, 
contained in steel cylinders, is an article of commerce. 

Although N 2 contains a smaller proportion of oxygen than 
any other oxide of nitrogen, it supports combustion almost as 
well as oxygen itself. A glowing splint bursts into flame when 
brought into N 2 0, just as it would do in 2 . 

557. Oxidation by Nitric Acid. — We are now ready to discuss 
oxidation by nitric acid. Since these reactions involve changes in 
electric charges on the nitrogen atom (501), let us go over the 
common reduction products of nitric acid with a view to assign- 
ing the electric charges to the nitrogen atoms present in them. 
In doing this two rules are followed; (1) oxygen in combination 
is double negatively charged except in the case of a peroxide; 



350 Introduction to General Chemistry 

and (2) hydrogen in combination is always single positively 
charged. Since the sum of the charges carried by the atoms in 
an electrically neutral substance must be zero, we may calculate 
the charge on one of the atoms of a compound if its charge is the 
only one not known. Table XXIII shows the probable charges 
on the nitrogen atom in the compounds we have studied. For 
convenience, charges greater than three in number are written 
with a superscript to show the number of charges. Thus N 5+ 
is equivalent to N +++++ . When nitric acid is reduced to 
N0 2 the original atom of nitrogen has gained one electron 
(N 5+ ->N 4+ ). To form N 2 it must gain five electrons, and to 
form NH 3 it must gain eight. Therefore it is a greater feat to 

TABLE XXIII 

Compounds Nitrogen Atoms 

HNO3, N 2 5 and nitrates N*+ 

N0 2 , N 2 4 N4+ 

HN0 2 , N 2 3 and nitrites N+++ 

NO N++ 

N 2 . . . N+ 

N 2 N° 

NH 3 . . N 

reduce nitric acid to ammonia than to nitric oxide. And in 
general we find that the more active the reducing agent the more 
we get of products like N 2 0, N 2 , and NH 3 . Thus the main 
products of the reaction of copper and nitric acid are NO and 
N0 2 (550) ; but when zinc reacts with nitric acid, N 2 0, N 2 , and 
NH 3 may be formed. The latter is found as NH 4 N0 3 , of course, 
if there is excess nitric acid available for combination. These 
results agree with our previous experience with these metals, 
since we have already found in the study of the electromotive 
series that zinc gives up electrons far more easily than does 
copper (491). 

The concentration of the solution, however, as well as the 
nature of the reducing agent, determines the final products of 
these reactions by reason of the reversible reaction 

3 N0 2 +H 2 0->2HN0 3 +NO. 



Nitric Acid and Oxides of Nitrogen 351 

We find, for example, that NO is the main product of the action 
of copper on dilute nitric acid, while with concentrated nitric 
acid N0 2 is chiefly formed. 

558. Reactions of Nitric Acid with Metals. — All the common 
metals except platinum and gold dissolve in nitric acid to form 
nitrates, though the latter are sometimes much hydrolyzed. It 
might be thought that zinc^would displace hydrogen from nitric 
acid, but as a matter of fact only metals in the electromotive 
series down to and including aluminum can generate hydrogen 
fast enough for it to escape the oxidizing action of nitric acid. 
The cases of copper and zinc already discussed may be taken as 
illustrative of these reactions. 

559. Reactions of Nitric Acid with Non-Metals. — Finely 
divided sulfur is oxidized to sulfuric acid by hot nitric acid. 
Iodine is oxidized to iodic acid, HI0 3 . Red phosphorus is 
oxidized to phosphoric acid (Care! see 591). Powdered charcoal 
(carbon) is oxidized to carbon dioxide. In all of these actions 
brown fumes are given off, showing the release of NO or N0 2 . 

560. Nitric Acid as a Chemical Solvent for Salts. — The fact 
that nitric acid is an oxidizing agent accounts for its ability to 
dissolve many salts on which hydrochloric acid fails to act. 
Thus the insoluble copper sulfide does not dissolve appreciably in 
HO (456) but does dissolve in dilute nitric acid, forming copper 
nitrate and sulfur, or sulfuric acid if the action with the acid is 
continued. So also mercurous chloride, insoluble in HC1, is 
dissolved easily by dilute HN0 3 to form mercuric nitrate. If 
neither of the ions of a salt can be oxidized nitric acid is no better 
chemical solvent than HC1. Thus neither acid can dissolve 
barium sulfate appreciably. 

561. Balancing of Oxidation and Reduction Equations. — 
Before leaving the subject of oxidation by nitric acid it is worth 
while to take up a systematic method for balancing oxidation 
and reduction equations, to be used when the proper coefficients 
cannot be easily arrived at by inspection. The following illus- 
tration of a useful method is given in detail. As has been said, 
copper reacts with concentrated nitric acid to give copper nitrate, 
nitrogen tetroxide, and water. We may first write down the 



35 2 Introduction to General Chemistry 

formulae of the substances taken and of the products in the form 
of an equation without the coefficients : 

Cu+HNO,->Cu(N0 3 ) 2 +N0 2 +H 2 0. 

Inspection of the formulae shows that the only atoms which 
change their valence are those of copper and those of nitrogen 
which go to form N0 2 (not those which go to form nitrate) . The 
equation for the exchange of charges by these atoms is as follows : 

Cu+ 2 N 5 +->Cu+ + + 2N<+ . 

The equation is balanced so that the number of electrons given 
up by the copper atoms equal those taken by the nitrogen. We 
may next write the equation to represent the change between the 
molecules which contain these atoms : 

Cu+2HN0 3 ->Cu++ + 2 N0 2 +H 2 0+0--. (i) 

On the right-hand side we need not trouble to assign positive 
ions to negative ions to form molecules unless so doing accounts 
for some of the actual products. N0 2 and H 2 are examples. 
We may next write equations to show the transformation of the 
Cu ++ and 0" into the products in which they finally appear, 

Cu++ + 2HN0 3 ->Cu(N0 3 ) 2 +2H+. (2) 

Gathering up the 0~~~ and H + from (1) and (2), we may com- 
bine them to form water, 

0-~+2H+->H 2 0. (3) 

Since all the products are accounted for, we may now secure the 
final equation by adding equations (1), (2), and (3) and canceling 
the terms which appear on both sides : 

Cu+ 4 HN0 3 ^Cu(N0 3 ) 2 +2N0 2 +2H 2 0. 

A little experience will show that there is a great similarity 
in oxidation and reduction equations, so that the detail which 
we have written out so laboriously will soon become so familiar 
that the intermediate equations need not be written out. How- 



Nitric Acid and Oxides of Nitrogen 353 

ever, until this facility is gained it is better to write out each step, 
as was done above. 

One more example, which is a difficult one, will now be taken 
up, namely the reaction between zinc and nitric acid to form 
zinc nitrate, ammonium nitrate, and water. We shall follow 
exactly the same scheme as before. The first step is to arrange 
the initial materials and end products in the form of an equation 
without the coefficients: 

Zn+HN0 3 ->Zn(N0 3 ) 2 +NH 4 N0 3 +H 2 0. 

Next we write and balance an equation to show the exchange in 
charges between the atoms which undergo oxidation and reduc- 
tion: 

4Zn+N 5 +-> 4 Zn+++N— " . 

The equation for the reaction of the molecules which contain 
these atoms is 

4 Zn+HN0 3 -> 4 Zn+++N + 3 0~+H+. (1) 

On the right-hand side of the equation we have not troubled to 
assign any of the positive ions to the negative ions, since in so 
doing we would not account for any of the actual products. 
Next we must account for the actual products one by one. 
These result from the union of the ions represented in (1) with 
the ions of the excess acid present: 

4Zn+++8HN0 3 ->4Zn(N0 3 ) 2 -f-8H+, (2) 

N— -f- 3 H+->NH 3 , (3) 

NH 3 +HN0 3 ->NH 4 N0 3 . (4) 

Gathering up the excess H + and 0~~ from equations (1), (2), 
(3), and (4) we may write 

3O— +6H+-» 3 H 2 0. (5) 

All products are now accounted for. We may add equations 
(1), (2), (3), (4), and (5) and cancel the terms which appear on 
both sides to secure the final equation 

4 Zn+ioHN0 3 ->4Zn(N0 3 ),+NH 4 N0 3 +3H,0. 



354 Introduction to General Chemistry 

562. Aqua Regia. — As has already been stated, gold is not 
attacked by nitric acid ; but this royal metal dissolves readily in 
the liquid made by mixing nitric acid with three times its volume 
of concentrated hydrochloric acid. Therefore the liquid solvent 
for the royal metal was called aqua regia (" royal water") by the 
alchemists of a thousand years ago. When aqua regia is gently 
warmed the liquid becomes yellow and gives off gases as the result 
of a reaction probably best represented by the equation 

3 HC1+HN0 3 ->2H£>+N0C1+CL. 

The substance NOC1 is called nitrosyl chloride. The action of 
aqua regia on gold converts the latter into auric chloride, AuCl 3 > 
a yellow salt easily soluble in water. Platinum, which like gold 
is also insoluble in either nitric or hydrochloric acid singly, dis- 
solves in aqua regia to form platinic chloride, PtCl 4 . 

563. Nitrosyl Chloride, NOC1. — This substance is a brownish- 
colored gas of very disagreeable odor. It is used extensively 
for bleaching flour. Less than one gram of NOC1 is required to 
bleach a barrel of flour. Nitrosyl chloride is readily acted on 

by water, thus : 

N0C1+H 2 0->HC1+HN0 2 . 

The nitrous acid formed is unstable and soon decomposes: 
3 HN0 2 ->HN0 3 + 2 NO+H 2 0. 

564. Nitrates. — Nearly every basic ion can form a nitrate. 
The nitrates are all easily, some extremely, soluble in water. A 
few nitrates derived from weak bases are hydrolyzed by water, 
forming difficultly soluble basic nitrates. Thus mercuric nitrate 
and water give Hg(N0 3 ) 2 *2HgO-H 2 0, called basic mercuric 
nitrate. Bismuth nitrate and water give the basic nitrate 
BiON0 3 , called bismuth subnitrate, used extensively in medicine. 

The nitrate ion, N0 3 ~, is colorless, and so also are all nitrates 
of colorless basic ions. 

565. Decomposition of Nitrates by Heat.— All nitrates are 
decomposed when they are heated to sufficiently high tempera- 
tures. They fall into three classes with respect to their behavior 
when heated. 



Nitric Acid and Oxides of Nitrogen 355 

1. When sodium nitrate, NaN0 3 , is heated, it first melts 
without decomposition to give a colorless liquid. At a higher 
temperature the liquid appears to boil. The gas given off is 
oxygen, formed as follows: 

2 NaN0 3 ->2NaN0 2 +0 2 . 

The other product, NaN0 2 , is sodium nitrite, the formation of 
which by the action of N0 2 on NaOH has already been mentioned 
(553)- Potassium nitrate, KN0 3 , when heated gives potassium 
nitrite, KN0 2 , and oxygen. 

2. The effect of heat on lead nitrate results in the following 

decomposition : 

2Pb(N0 3 ) 2 ->2PbO+ 4 N0 2 +0 2 . 

This reaction is typical of the behavior of the nitrates of most 
metals excepting those of the alkali group (Na, K, etc.). In 
some cases the oxide of this metal is also decomposed by heat, 
so that the free metal is formed. This is the case with the 
nitrates of mercury and silver. 

3. The action of gentle heat on ammonium nitrate, which 
takes place thus, 

NH 4 N0 3 ->N 2 0+ 2 H 2 0, 

has already been described (556) . 

566. The Equilibrium between Nitrogen and Oxygen. — We 

are now ready to take up in detail the consideration of the prob- 
lem of the manufacture of nitric acid from atmospheric nitrogen. 
We shall begin by the study of the reversible reaction 

N 2 +0 2 ^2NO. 

This reaction takes place so slowly below 8oo° (a bright-red heat) 
that ordinary observation would lead one to conclude that 
nitrogen and oxygen have no tendency to unite, and that nitric 
oxide has no tendency to dissociate into its constituents. At 
very high temperatures each reaction takes place more rapidly, 
so that at the very high temperature of 3000 equilibrium is 
reached in a small fraction of a second. At 3000 only about 
5 per cent of the N 2 and 2 are combined as NO in the equilibrium 



356 



Introduction to General Chemistry 



mixture. At lower temperatures the equilibrium mixture con- 
tains still smaller proportions of NO. Table XXIV shows the 
percentage by volume of NO in the equilibrium mixture at each 
of several temperatures, and also the length of time required to 
reach practically complete equilibrium (that is, within one-tenth 
of i per cent complete). Inspection of the table shows that the 
higher the temperature the greater the proportion of NO formed 
and the more quickly the attainment of equilibrium. 





TABLE XXIV 




Temperature 


Per Cent of NO 


Time 


I sOO° 


O.I 
1.2 
2.6 

5-3 


12 days 
30 sec. 
. 1 sec. 


2000° 


2 sOO° 


° o 

^OOO 


. 0004 sec. 





567. The Manufacture of Nitric Acid from Air. — It is a 
simple matter to make nitric acid from nitric oxide, NO. We 
have learned that NO unites readily at ordinary temperatures 
with oxygen of the air to form nitrogen tetroxide, 

2 N0^0 2 +2N0 2 . 

The latter gas reacts with water to give nitric acid and nitric 

oxide, 

3 N0 2 +H 2 0-> 2 HN0 3 +NO; 

and since the NO by direct union with oxygen of the air passes 
readily into N0 2 , the whole of the NO can finally be converted 
into nitric acid. The reactions discussed in this and the fore- 
going paragraphs are now carried out on a very large scale for 
the manufacture of nitric acid from atmospheric nitrogen. The 
difficult part of the process is the preparation of NO. We see 
from the table that the higher the temperature to which the 
mixture of N 2 and 2 is heated the greater the proportion of 
NO present in the equilibrium mixture. Now it is not at all 
difficult to heat air, which is simply a mixture of N 2 and 2 , to 
3000 or even higher by means of an electric arc. The air at 
this temperature then contains 5 per cent or more of NO, which 



Nitric Acid and Oxides of Nitrogen 



357 



is a very satisfactory proportion. But a most serious difficulty 
is now encountered : this gas mixture is at a dazzling white heat, 
3000 ! It must be cooled almost to ordinary temperatures 
before it will unite with more oxygen to form N0 2 ; and the latter 
must be nearly cold before it is combined with H 2 to produce 
HNO3. Since the cool- 
ing cannot be accom- 
plished instantaneously, 
the temperature for an 
appreciable time will be 
between 25oo°and 2000 . 
Therefore the reaction 

N 2 +0 2 ±5 2NO 

will reverse to a greater 

or less extent, since at 

2500 only 2.6 per cent 

of NO is stable and at 

2000 only 1 . 2 per cent. 

If the mixture which 

contained 5 per cent of 

NO at 3000 remains at 

2000 for 30 seconds only 

1 . 2 per cent of NO will 

be left! This makes it 

imperative to cool the gas mixture very rapidly to a temperature 

below 1000 , where the speed of decomposition of NO becomes 

negligible. 

One of the most successful practical methods of making NO 
will now be described. 

568. The Birkland and Eyde Process. — This process was 
invented in 1903 by two Norwegians, Birkland and Eyde. Air 
is passed through an immense electric arc and then cooled as 
rapidly as possible. The mixture of unchanged air and nitric 
oxide is then used to produce nitric acid by means of the reactions 
already discussed (567). The Birkland-Eyde electric furnace 
is shown in cross-section in Fig. 90. An arc formed at A between 




Fig. 90 



358 



Introduction to General Chemistry 



the electrodes which enter from the front and back is spread out 
by the action of a powerful magnet N-S into a great electric 
flame which fills the space B-B. The wires of the magnet are 
shown at M-M. The arc is inclosed by refractory walls, pierced 
by numerous holes. (A refractory substance is one which is 
incombustible and may be heated to a high temperature.) Air 
passes through the furnace in the manner shown by the arrows. 
The temperature of the perforated walls inclosing the arc does 
not exceed 8oo°, owing to the cooling effect of the inflowing air. 
In the arc the gases reach a temperature of 3500 , but are sub- 
sequently quickly cooled by the walls, so that they escape at 




^ 



+ ^ 




Fig. 91 



about 1000 . The NO content is about 2 per cent. In addition 
to the furnace just described, several others accomplishing 
similar results are in practical use. 

569. Conversion of Nitric Oxide into Nitric Acid. — Figure 91 
illustrates diagrammatically a typical plant for the manufacture 
of nitric acid from air. 

A pump, A (air compressor), delivers air to the electric 
furnace, F, from which the escaping mixture of air and NO at 
1000 passes into a cooler, C, where the temperature is lowered 
to 500 . The gases are now passed under steam boilers, B, to 
generate steam for use in the plant, and are thereby still further 
cooled. The reaction 

2 NO+0 2 ->2N0 2 



now takes place at about 50 in the vessel, D, there being a large 
excess of 2 still present. The absorption tower, T, into the 



Nitric Acid and Oxides of Nitrogen 359 

base of which the gases now pass, is filled with pieces of quartz 

and is supplied with a stream of water at the top. Here the 

reaction 

3 N0 2 +H 2 0->2HN0 3 +NO 

takes place. Nitric acid is drawn off at the base of the tower, 
while the NO and -air pass out at the top. The escaping NO is 
not lost, since it can combine with more 2 to form N0 2 , which is 
also converted into nitric acid, so that finally all NO is changed 
into HNO3. Large amounts of nitric acid are now manufactured 
in this way. 

570. The Manufacture of Nitric Acid from Ammonia. — The 

conversion of NH 3 into HN0 3 was mentioned in the foregoing 

chapter. When a mixture of one volume of NH 3 and two 

volumes of oxygen or ten volumes of air is passed over a gauze 

of fine platinum wire at about 700 the following reaction takes 

place : 

NH 3 + 2 2 ->HN0 3 +H 2 0. 

The platinum acts catalytically, and the speed of the reaction is 
so great that contact with the gauze for o. 01 second is sufficient. 
Considerable heat is given out, so that by regulation of the rate 
of flow of the gases the most favorable temperature is maintained. 
Although this reaction has been known since 1830, it did not 
assume commercial importance until after cheap, synthetic 
ammonia became available. Since the beginning of the war a 
large proportion of the nitric acid required by Germany has been 
made from NH 3 by this process. As stated in the foregoing 
chapter (526), an immense amount of nitric acid will be made 
at the new American works at Muscle Shoals, Alabama, from 
NH 3 made by the cyanamid process. 

571. Uses of Nitric Acid. Explosives. — With very few ex- 
ceptions all practical explosives are made by the use of nitrates, 
or nitric acid. The oldest explosive is gunpowder (now better 
known as black powder). This is a mechanical mixture of 
charcoal, sulfur, and potassium nitrate. Modern explosives are 
made by the action of nitric acid on various substances, such as 
glycerine, cotton, phenol, toluene, etc. The chemistry of these 



360 Introduction to General Chemistry 

nitrated substances, all of which are carbon compounds, will be 
considered in chapter xxvi. Suffice it to say here that without 
nitric acid we should have no modern explosives, and without 
modern explosives present forms of warfare would not exist. 
We should also remember, however, that the arts of peace are 
quite as dependent upon explosives as are those of war. Modern 
methods of mining and quarrying would be impossible without 
explosives; and without the aid of these marvelous chemical 
products the Panama Canal could not have been dug. 

Nitric acid is an indispensable reagent in the manufacture 
of most artificial dyes. It also has many other important uses. 

572. Gunpowder, Black Powder. — The history of the dis- 
covery of gunpowder, usually ascribed to Roger Bacon, is rather 
obscure. Gunpowder seems first to have been used in cannon 
toward the end of the thirteenth century. This explosive is 
a black, granular substance, composed of about 74 per cent 
potassium nitrate, 16 per cent charcoal, and 10 per cent sulfur. 
In making gunpowder the finely powdered components are 
moistened with water, and the mass is thoroughly mixed. The 
product is pressed into cakes; these are dried and broken into 
grains of suitable size. The explosion of gunpowder results in 
the change of KN0 3 , S, and C into C0 2 , S0 2 , N 2 , and K 2 S. Of 
the products all but the potassium sulfide are gases. The very 
large volume of gases formed in a fraction of a second is the cause 
of the result which we call an explosion. These gases, at the 
instant of their liberation, are at a very high temperature and 
occupy only the same space as that of the powder. They there- 
fore exert an enormous pressure on all sides. In a gun the pro- 
jectile is shot out, while the gun itself recoils. 

573. Ammonium Nitrate as an Explosive. — Ammonium 
nitrate, NH 4 N0 3 , is a very important explosive. We have seen 
that when gently heated it decomposes, thus : 

NH 4 N0 3 ->N 2 0+2H 2 0; 

but when heated to a very high temperature or when exploded 
by a powerful detonator the decomposition is more complete: 

2 NH 4 N0 3 ->2N 2 + 4 H 2 0+0 2 . 



Nitric Acid and Oxides of Nitrogen 361 

At the high temperature of the explosion H 2 is of course gaseous. 
A detonator is a substance that explodes readily and produces a 
sharp shock that starts the explosion of a large quantity of 
another less easily exploded substance. Ammonium nitrate is 
not easily exploded. It may therefore safely be used in shells 
fired from cannon without risk of its premature explosion. A 
time or impact detonator is required for such shells. Enormous 
quantities of NH 4 N0 3 were used in the war. 

574. The Cycle of Nitrogen in Nature. — We have already 
learned (511) that the flesh of all animals is made up of com- 
pounds of nitrogen (with carbon, hydrogen, and oxygen chiefly). 
Animals cannot derive their required nitrogen from the air or 
from any of the simple nitrogen compounds so far studied. 
Much more complex nitrogen compounds (proteins) are required 
for the food of animals (including man). These nitrogenous 
animal foods are obtained from the flesh of other animals and 
from plants, particularly cereals. Plants have the power of 
building up proteins from simple nitrogen compounds, such as 
ammonium salts and nitrates, which are present in the soil. A 
small group of plants (especially the legumes: clover, alfalfa, 
etc.) can, by virtue of bacteria that infest their roots, take up 
free nitrogen and convert it into proteins. All other plants 
thrive only on soils containing compounds of nitrogen. Soils 
deficient in combined nitrogen are greatly improved in fertility 
by the application of ammonium salts or of nitrates as fertilizers. 
In normal times about three-fourths of all ammonium salts and 
nitrates used in America are employed as fertilizers. A large 
part of the nitrogen of animal food is eliminated in the form of a 
rather simple substance, urea, CO(NH 2 ) 2 . This substance 
changes slowly into C0 2 and NH 3 by the action of water, 

CO(NH 2 ) 2 +H 2 0->C0 2 + 2 NH 3 . 

Urea itself, as well as NH 3 , can serve as plant food. 

The decay of animal remains and refuse results in the change 
of their proteins into simpler substances. Ammonium salts are 
first formed and later are oxidized to nitrates. Thus nitrogen 
in nature passes through an endless cycle of changes. 



362 Introduction to General Chemistry 

575. The Problem of the Fixation of Nitrogen. — In this 
chapter and the foregoing chapters we have tried to point out 
the practical importance of the compounds of nitrogen. The 
abundance of food is determined by the fertility of the soil. 
Fertility may be conserved and increased by means of nitrogenous 
fertilizers. Aside from animal refuse and manures and the 
ammonia obtained as a by-product of the manufacture of gas 
and coke, the great deposits of sodium nitrate in Chile were 
until recently the* only important source of nitrogenous fer- 
tilizers. These deposits are by no means inexhaustible. They 
may last one hundred years at the present rate of consumption. 
Fortunately, within the last twenty-five years several methods 
have been developed for making nitrogen compounds directly 
from the free nitrogen of the air. All such methods are now 
referred to as processes for the fixation of nitrogen. We have 
at present three successful methods, each of which is in extensive 
use. These are: (1) the making of nitric acid from air (566), 
(2) the making of cyanamid from calcium carbide and atmos- 
pheric nitrogen (526), and (3) the synthesis of ammonia from 
hydrogen and atmospheric nitrogen (520). We have also seen 
that ammonia is readily made from cyanamid, and that nitric 
acid is easily produced by the oxidation of ammonia by air. 
Since the nitrogen contained in the air over each square mile of 
the earth's surface amounts to nearly 20,000,000 tons, the supply 
may be said to be inexhaustible. 

The final question to be settled before deciding upon the 
practicability of any technical process is whether it is eco- 
nomically sound; in other words, will it pay? With regard to 
the three general methods of fixing nitrogen above mentioned, 
it may be said that the direct manufacture of nitric acid from 
air requires exceedingly cheap electric power. The other two 
general methods can be used where power is more expensive. 
At present all three processes are profitable. Only time will 
decide which will prove most economical. In any case we may 
feel satisfied that the problem of the fixation of nitrogen has been 
solved. 



CHAPTER XXIII 
PHOSPHORUS 

576. Review. — We have already had some experience with 
the chemistry of phosphorus and its compounds. The burning 
of the free element was used to show the presence of oxygen in 
air (10). In chapter ix phosphoric acid and some of its salts 
were described (158, 159). The importance of phosphates for 
the life of plants and animals was pointed out (160), and their 
widespread distribution in nature was mentioned. One method 
of preparing free phosphoric acid was also given. In chapter xii 
(247) the chemistry of phosphorus was continued in the descrip- 
tion of the union of phosphorus with the halogens to form 
trivalent and pentavalent phosphorous compounds, and the 
reactions of these compounds with water to set free the hydrogen 
halides and an oxygen acid of phosphorus, as illustrated by the 
following typical equations: 

PC1 3 + 3 H 2 0->H 3 P0 3 + 3 HC1, 
PBr s +4H 2 0->H 3 P0 4 +5HCl. 

In chapter xxii the powerful dehydrating property of the pen- 
toxide of phosphorus was mentioned in the formation of the 
anhydride of nitric acid : 

2HN0 3 +P 2 O s ^2HP0 3 +N 2 O s . (555) 

We shall now take up a systematic study of the element 
phosphorus and its compounds. 

577. The Discovery of the Element. — The free element, 
which does not occur in nature, was discovered by Brandt, of 
Hamburg, sometime before 1669. In searching for a suitable 
solvent for turning silver into gold he was led to ignite a mixture 
of dried urine and charcoal in a clay retort. Since the urine 
contained phosphates he was really carrying out one of the 

3 6 S 



364 Introduction to General Chemistry 

methods now in use for preparing free phosphorus. When we 
have studied some of the peculiarities of the free element we shall 
realize what a stir this new substance must have made in the 
thinking world of that day. Historians record with unction 
that Kraft, who purchased the secret of its preparation from 
Brandt, exhibited it before many of the crowned heads of Europe. 

578. The Physical Properties of White Phosphorus. — At 
ordinary temperatures the free element in pure form is a white, 
waxy, translucent solid of specific gravity 1.8. It does not con- 
duct electricity. It is usually kept under water, since when it 
is exposed to air it soon takes fire. At ordinary temperatures 
the solid is slightly volatile. Calculations made from the density 
of the vapor show that it has a formula of P 4 at temperatures up 
to 1,500°. It melts at 45° and boils at 287°. It is not soluble in 
water but is readily soluble in many solvents, such as carbon 
disulfide, turpentine, olive oil, and many of its own compounds. 
The so-called yellow phosphorus of commerce is white phos- 
phorus made yellow by the presence of impurities. 

579. The Conversion of White Phosphorus into Red Phos- 
phorus. — When white phosphorus is exposed to light it darkens 
in color, owing to the formation of a new form called from its 
color red phosphorus. Samples of white phosphorus which have 
been kept for some time show a coating of this substance. At 
higher temperatures the change from white to red goes on more 
rapidly. It can be hastened also by the presence of a trace 
of iodine, which acts as a catalyzer. 

580. The Manufacture of Phosphorus. — Though some phos- 
phorus is manufactured from bone ash, most of it is prepared 
from phosphate minerals. The principal materials are phos- 
phate rock and apatite. The first varies in composition from 
Ca 3 (P0 4 ) 2 to the composition of apatite, Ca 3 (P0 4 ) 2 CaF 2 , or 
Ca 3 (P0 4 ) 2 CaCl 2 . 

The most widely used method of making phosphorus con- 
sists in heating a mixture of sand, phosphate rock, and charcoal 
to high temperatures in an electric furnace, so designed that the 
phosphorous vapor distils off and condenses under water. The 
residual materials of the reaction are withdrawn from time to 



Phosphorus 365 

time, and new mixtures are put in in a continuous process. The 
equation of the reaction follows : 

2Ca 3 (P0 4 ) 2 +6Si0 2 +ioC->6CaSi0 3 +ioCO+P 4 . 

Red phosphorus is made on a commercial scale by heating 
yellow phosphorus in the absence of air. 

581. The Properties of Red Phosphorus. — Although red 
phosphorus is often given the title amorphous phosphorus, which 
means that it is without crystalline form, it is really a crystalline 
powder.. Its specific gravity is 2.1. Commercial red phos- 
phorus often contains among other impurities small traces of 
yellow phosphorus and some phosphoric acid. It is on this 
account sticky and hygroscopic (water-absorbing), although if 
free from these impurities it is a dry powder quite insoluble in 
water and also in most other solvents. It has far greater stability 
than white phosphorus, as is evidenced by the fact that it must 
be heated to 240 before it will take fire in air, and that it may 
be kept in contact with the latter at ordinary temperatures for 
years without alteration. If pure it is non-poisonous. It is 
for these reasons to be preferred to white phosphorus in all 
laboratory experiments in which the substitution can be made. 
The chemistry of red phosphorus parallels that of yellow phos- 
phorus, but in every case the reactions are less violent. 

582. Allotropic Forms. — White and red phosphorus are called 
allotropic forms of phosphorus. The noun allotropy translated 
means simply "other modes." We shall find many examples of 
elements which exist in more than one form. We have already 
met ordinary oxygen and ozone (316), allotropic forms of the 
same element. We found that these two forms differed from 
each other in the number of atoms and in the energy content of 
the molecule. The same causes undoubtedly account for the 
differences in the two forms of phosphorus. 

583. The Slow Oxidation of Phosphorus. — If white phos- 
phorus is exposed to moist air the heat given out by its oxidation 
will slowly raise its temperature to the melting-point, and it then 
takes fire. It is this fact that makes the storage of white 
phosphorus under water necessary. 



366 Introduction to General Chemistry 

If a stick of moist phosphorus is viewed in the dark it is seen 
to give out a soft, yellow light. It is from this property that 
the element received its name, which means " bearer of light." 
We can imagine what an interest the discovery of this substance 
aroused in a superstitious age ! The appearance of this light is 
apparently due to the slow oxidation which is in progress; for if 
conditions are such that oxidation cannot occur, no light is 
given out by the phosphorus. The light is simply a part of the 
energy of the reaction, which appears in this unusual form 
instead of as heat. 

584. Phosphorescence. — When light is given off by bodies at 
ordinary temperatures they are said to phosphoresce, and the 
process is called phosphorescence. But the same phenomenon 
is shown by substances other than phosphorus. Moist, decaying 
wood often glows in the dark as the result of the liberation of 
energy in the form of light. The so-called fox fire seen in 
forests after night is an example. We have another case in the 
process by which the firefly gives out light from parts of the 
surface of its body. All our processes of making light are very 
wasteful, since high temperatures must be maintained, and so 
most of the energy used must go into heat, and only a small 
portion is changed into light. Hence the process by which the 
firefly gives out light at the temperature of its body is indeed 
marvelous, and its imitation has been the despair of scientists 
for generations. 

585. Danger in the Use of White Phosphorus. — In conclusion 
it should be especially emphasized that white phosphorus is a 
very dangerous reagent. // should never be touched by the hands 
directly, since it sticks to the skin and the heat of the hand is 
sufficient to ignite it. 

586. Matches. — On account of the serious harm to workers 
with yellow phosphorus, by reason of poisonous vapor, the mak- 
ing of yellow-phosphorus matches has been made illegal in all 
the great countries of the world. These. matches were at one 
time very popular because of their easy ignition. Matches are 
now made with red phosphorus, phosphorus sulfide, P 4 S 3 , or 
other non-poisonous compounds of phosphorus. The ordinary 



Phosphorus 367 

match is a stick of non-resinous wood dipped first in paraffin 
and then in a preparation to make the ignition head. The latter 
is usually a mixture of a good oxidizing material like lead dioxide 
or potassium chlorate and a reducing agent such as red phos- 
phorus or phosphorus sulfide. With these is mixed some 
powdered glass to increase the friction when the match is struck. 
The material of the head is kept in place by glue or varnish. 
Friction starts the action between the oxidizer and the reducer, 
and the heat of this reaction kindles the paraffin, which in turn 
kindles the wood. The heads of safety matches are made of less 
easily reacting mixtures, usually sulfur and potassium chlorate. 
They are ignited by striking on a prepared surface coated 
with red phosphorus, antimony trisulfide, and glue. Match 
sticks are often treated with phosphoric acid or sodium phosphate 
to prevent " afterglow," the formation of glowing coals of char- 
coal after the flame of the match has been blown out. 

587. Smoke Screens. — A new use for phosphorus was 
developed during the war. It was discovered that the most 
effective smoke screens for concealing troop movements were 
produced by means of burning phosphorus. The explosion of 
shells filled with phosphorus ignited the latter, which burned to 
give dense white clouds of the pentoxide, P 2 5 . Large quantities 
of phosphorus were used for this purpose. 

588. Compounds of Phosphorus. — Phosphorus forms com- 
pounds with many elements. With hydrogen it unites to form 
several hydrides. Of these, phosphine, PH 3 , has been of particu- 
lar interest, since it apparently resembles ammonia. It is a 
poisonous, ill-smelling gas, little soluble in water, but like am- 
monia it unites with acids to form salts, though far less readily. 
For instance, phosphine and hydrogen chloride unite at low 
temperature to form a cloud of solid phosphonium chloride: 

PH 3 +HCtePH 4 Cl. 

Some of the halogen compounds of phosphorus have been 
already studied. These are very important substances used to 
synthesize many valuable halogen compounds. The formation 
of hydrogen bromide or hydrogen iodide by the action of water 



368 Introduction to General Chemistry 

on the tribromide or the tri-iodide is typical of these important 
reactions and should be reviewed (256, 264). 

With oxygen, phosphorus unites to form P 2 3 , P 2 4 , and 
P 2 O s , respectively. The first two are formed when phosphorus 
burns in a poor supply of oxygen; the last is formed when the 
burning takes place in an excess of oxygen. All three are white 
hygroscopic solids. 

Phosphorus trioxide unites with cold water to form phos- 
phorous acid: 

3 H 2 0+P 2 3 ^ 2 H 3 P0 3 . 

The latter is a very unstable acid. When the solution is 
heated some of the molecules of the acid are oxidized to phos- 
phoric acid at the expense of others, which are reduced, forming 
phosphine. 

Phosphorus pentoxide unites with water with great vigor, 
forming first metaphosphoric acid, HP0 3 . The stability of this 
acid and the completeness of the reaction make phosphorus 
pentoxide our most effective drying agent. The last traces of 
moisture are taken from air exposed to this oxide. 

589. The Phosphoric Acids. — When phosphorus pentoxide 
is thrown into water spattering is likely to occur because of the 
heat given out in the formation of metaphosphoric acid : 

H 2 0+P 2 O s ^2HP0 3 . 

The excess water may be evaporated, leaving the pure acid, a 
colorless, glassy solid. It changes to orthophosphoric acid, 
H 3 P0 4 , if it is kept a long time in water solution. The change 
may be carried out in the course of an hour if the metaphos- 
phoric acid solution is mixed with a strong acid and heated on 

the water bath: 

HP0 3 +H 2 O^H 3 P0 4 . 

The action is hastened by the catalytic action of the hydrogen 
ion of the strong acid. 

590. The Properties of Orthophosphoric Acid. — Orthophos- 
phoric acid, usually called simply phosphoric acid, forms color- 
less, odorless crystals, which are deliquescent and very soluble 



Phosphorus 369 

in water. The acid is usually sold in the form of a concentrated 
solution of sirupy consistency. Its dilute solution has an agree- 
able sour taste and is not poisonous. This acid cannot be 
oxidized and is very stable toward most reducing agents. It is 
a moderately strong acid (409) ; but since it is tribasic (159) we 
shall consider, in some detail, in a separate section, its interesting 
mode of ionization. 

Orthophosphoric acid is not volatile. For this reason it 
reacts at elevated temperatures with salts of volatile acids like 
chlorides, bromides, and nitrates to form phosphates and to set 
free the volatile acids. For example, 

H 3 P0 4 +NaBr->NaH 2 P0 4 +HBr. 

When phosphoric acid is heated to 2io°-2i5° in an open vessel 
for about an hour it is converted, by slow loss of water, into 
pyrophosphoric acid, H 4 P 2 7 : 

2 H 3 P0 4 ->H 4 P 2 7 +H 2 0. 

If the heating is continued longer and the temperature is raised 
to 400 or higher, the product slowly loses more water, giving 
finally metaphosphoric acid, HP0 3 : 

H 4 P 2 7 -> 2 HP0 3 +H 2 0. 

Metaphosphoric acid volatilizes slowly at a red heat (7oo°-8oo°) . 

591. The Preparation of Phosphoric Acid. — Phosphoric acid 
is easily made in the laboratory by boiling red phosphorus with 
concentrated nitric acid. (Caution ! White phosphorus must 
not be used.) The violent reaction produces red fumes of 
nitrogen tetroxide, N0 2 , and gives finally a colorless, sirupy 
solution, from which the excess of HN0 3 can be removed by 
heating, since HN0 3 is easily volatile, while H 3 P0 4 is not. 
Incomplete oxidation gives more or less phosphorous acid, 
H 3 P0 3 . The latter is readily oxidized to H 3 P0 4 by sufficient 
boiling nitric acid. 

Phosphoric acid is made commercially by the double decom- 
position of calcium phosphate and sulfuric acid (158) and sub- 
sequent purification of the product. The crude acid is very 



37° Introduction to General Chemistry 

likely to contain magnesium and calcium salts and also some 
arsenic acid, H 3 As0 4 , the first coming from the rock used and 
the second mostly from the acid employed, though sometimes 
from the rock as well. 

592. Ionization of Phosphoric Acid. — According to the ionic 
theory phosphoric acid dissociates as follows : 

H 3 P0 4 ^H++H 2 P0 4 -, 
H 2 P0 4 -^H++HP0 4 — , 
HP0 4 -~^H++P0 4 . 

In a water solution equilibrium is established between the three 
pairs of opposed reactions, and each equilibrium is dependent 
on the other two. In all cases of this kind the first dissociation 
goes on to a greater degree than the second, the second to a 
greater degree than the third, etc. In a solution of phosphoric 
acid containing one-tenth formula weight per liter at 18 the 
total concentration of hydrogen ion is 0.0282 and that of the 
primary phosphate ion, H 2 P0 4 ~, virtually the same; the con- 
centration of the secondary phosphate ion, H 2 P0 4 ~~, is 
calculated to 0.0000002, while that of the phosphate ion, 
P0 4 , is very much smaller. 

It is important that we understand that, inasmuch as H 2 P0 4 ~ 
and HP0 4 both dissociate to give hydrogen ion, they are acids, 
just as is phosphoric acid itself. Since, however, their degree of 
dissociation is far less than that of phosphoric acid, even at 
greater dilutions they are far weaker acids. Obviously HP0 4 ~~ 
is much weaker than H 2 P0 4 ~. These acids are furnished by 
the soluble and highly ionized salts NaH 2 P0 4 and Na 2 HP0 4 as 
a result of their ionization: 

NaH 2 P0 4 ^Na++H 2 P0 4 ", 
Na 2 HP0 4 ^ 2 Na++HP0 4 — . 

593. Phosphate Baking Powder. — The ion H 2 P0 4 ~ is a weak 
acid of about the same strength as carbonic acid, but it has an 
advantage over the latter in being more stable in water solution. 
As a result, when water is added to a mixture of sodium hydrogen 
carbonate, NaHC0 3 , and sodium acid phosphate, NaH 2 P0 4 , we 



Phosphorus 371 

get, as we expect, an effervescence due to the escape of carbon 
dioxide from the solution of carbonic acid formed. Thanks to 
the escape of the gas the action becomes complete : 

NaHC0 3 +NaH 2 P0 4 ^Na 2 HP0 4 +H 2 0+C0 2 . 

This is the reaction made use of in phosphate baking powders. 
The essential ingredients are sodium bicarbonate and an acid 
phosphate which may be the sodium, calcium, or magnesium 
salt, or mixtures of them, together with some inert material, 
such as starch, which serves as a dilutent. When water is added 
to the mixture of the dry salts and flour the tiny bubbles of 
carbon dioxide given off make the dough spongy. 

594. The Phosphates. The Sodium Salts. — Primary sodium 
phosphate, or sodium acid phosphate, NaH 2 P0 4 , is prepared in 
solution by adding one molecular weight of sodium hydroxide 
for every molecular weight of phosphoric acid : 

H 3 P0 4 +NaOH->NaH 2 P0 4 +H 2 0. 

That NaH 2 P0 4 alone forms, and not a mixture of the three 
possible phosphates, we can understand if we remember that the 
secondary and tertiary salts are salts of weaker acids than 
H 3 P0 4 ; and so if momentarily formed they would undergo 
double decomposition with unused phosphoric acid to form 
NaH 2 P0 4 . A solution of NaH 2 P0 4 is acid to litmus, as we would 
expect from the presence of the acid H 2 P0 4 ~. 

Common sodium phosphate, Na 2 HP0 4 , the secondary salt, 
is prepared by adding two molecular weights of sodium hydroxide 
for every molecular weight of phosphoric acid in solution. Solu- 
tions of secondary sodium phosphate are alkaline to litmus, 
owing to a slight hydrolysis (436). 

HP0 4 ~~ is so weak an acid that when three molecules of 
sodium hydroxide are furnished for every molecule of phosphoric 
acid in solution the action is far from complete, and we have only 
a small part of the possible Na 3 P0 4 formed when equilibrium is 
established. However, we can force the reaction to yield the 
desired salt by evaporating off the water, or better by furnishing 



37 2 Introduction to General Chemistry 

a very large excess of sodium hydroxide and then evaporating 
the water. In either case the tertiary or normal sodium phos- 
phate will separate from solution. As we would expect, if we 
redissolve this salt we get back the same solution made originally, 
and we find that as a result of the hydrolysis involved the solution 
is strongly alkaline. 

As solids, all three salts are white crystalline substances and 
are hydrated, as indicated by the following formula: 

NaH 2 P0 4 . 4 H 2 0; Na 2 HP0 4 • 1 2H 2 ; Na 3 P0 4 -i2H 2 0. 

When heated the primary and secondary salts both give off 
water and form sodium metaphosphate and pyrophosphate 
respectively. The tertiary phosphate remains unchanged: 

NaH 2 P0 4 ->NaP0 3 +H 2 0, 
2 Na 2 HP0 4 ->Na 4 P 2 7 +H 2 0. 

595. The Titration of Phosphoric Acid, Normal Solution. — 
If phenolphthalein is used as the indicator and phosphoric acid 
is titrated with sodium hydroxide (137), the color change of the 
indicator will be observed when two formula weights of sodium 
hydroxide have been added for every formula weight of the acid. 
This is the point at which the acid has been changed to the 
secondary salt, Na 2 HP0 4 . Therefore a liter of a solution of 
phosphoric acid which contains one-half formula weight of the 
acid will neutralize one liter of a normal sodium hydroxide solu- 
tion and has the same neutralizing power as have solutions of 
hydrochloric acid or nitric acid which contain one formula 
weight of the acid per liter respectively, or of sulfuric acid which 
contain one-half formula weight of the acid per liter. A half 
molar solution of phosphoric acid is therefore a normal solution 
of the latter. 

596. The Bead Test. — Sodium metaphosphate is much used 
in analytical chemistry for the so-called metaphosphate bead 
tests. The bead tests are carried out by first fusing micro- 
cosmic salt, sodium ammonium hydrogen phosphate, on the 
end of a platinum wire. This salt, which is used because it 



Phosphorus 373 

is easily prepared in pure form, decomposes, giving sodium 
metaphosphate as follows: 

NaNH 4 HP0 4 ->NaP0 3 +H 2 0+NH 3 . 

Then a tiny speck of the material to be tested is picked up on the 
bead and fused into it. Salts of metals fuse into the bead, 
forming phosphates, which often are characteristically colored 
and thus indicate the metal present in the original substance. 
Presumably the salt of the metal first forms an oxide and then 
unites with the metaphosphate to form an orthophosphate, a 
reaction which is really a reverse of the formation of the meta- 
phosphate from the orthophosphate described above. For 
instance, copper oxide would be dissolved, forming sodium 
copper phosphate, thus: 

CuO+NaP0 3 ->NaCuP0 4 . 

597. Qualitative Tests for Phosphates. — Phosphates are 
usually tested for by the formation of yellow ammonium phos- 
phomolybdate, (NH 4 ) 3 P04iiMo0 3 6H 2 0, which is insoluble in 
nitric acid solution. As the formula indicates, the successful 
operation of this test necessitates an enormous excess of the 
reagent ammonium molybdate. 

. Owing to the fact that solutions of the metaphosphates and 
pyrophosphates are stable and only form the orthophosphate 
after the lapse of considerable time, unless they are heated with 
strong acids, it is desirable to be able to distinguish between 
the three kinds of phosphates. This is done on the basis of the 
fact that silver orthophosphate, Ag 3 P0 4 , is insoluble in water 
and yellow, whereas the other silver phosphates, AgP0 3 and 
Ag 4 P 2 7 , though also insoluble, are white. The metaphosphate 
is distinguished from the pyrophosphate by the fact that the 
free acid which can be made from the salt will coagulate albumen, 
commonly known in the form of white of egg, while pyrophos- 
phoric acid will not. As a matter of fact this test distinguishes 
the metaphosphoric acid from the or triphosphoric acid as well. 

598. Use and Production of Phosphates. — Phosphates and 
phosphoric acid are used extensively in medicine, but the great 



374 Introduction to General Chemistry 

bulk of the phosphate production supplies the match industry 
and the still greater demands of agriculture. The meeting of 
the latter need is of incalculable importance, since the phos- 
phate taken from the land by the crops must be put back or the 
soil will become sterile. The natural phosphates are mined on 
an enormous scale either for direct use as fertilizer or for the 
production of superphosphate (160), to be used for the same 
purpose. 



CHAPTER XXIV 
SULFUR AND ITS COMPOUNDS 

599. Introduction. — Many isolated cases of the use of sulfur 
compounds have already been encountered. Some of the 
properties of sulfuric acid were given (93), and in subsequent 
work its use was described in the preparation of hydrochloric 
(103), nitric (104), and phosphoric (158) acids, and also in the 
making of many sulfates and acid salts. Hydrogen sulfide was 
used as 'a reducing agent (339), the most important instance 
being the preparation of hydrogen iodide by the action of 
aqueous hydrogen sulfide upon iodine. Sulfurous acid was also 
shown to be a reducing agent because of its easy oxidation to 
sulfuric acid (340) . As a matter of fact, so commonly are these 
and other sulfur reagents used by chemists, because of their 
cheapness and wide range of adaptability, that they may be 
counted among the most important tools of the trade. It is 
therefore from the point of view of possessing ourselves of a 
working knowledge of these important reagents that we are 
next to undertake the further study of sulfur and its com- 
pounds. 

600. The Physical Properties of Sulfur. — The commonest 
form of free sulfur is called rhombic sulfur from the shape of its 
crystals. It is light yellow in color, not soluble in water, slightly 
soluble in alcohol or ether, but very soluble in some of its own 
compounds, such as carbon disulfide, CS 2 (546). Another form 
of sulfur, monoclinic sulfur, which is also named from the shape 
of its crystals, has the same color as rhombic sulfur and is 
soluble in the same solvents but to a different degree. At 95 , 
the so-called transition point, these two forms of sulfur can exist 
in contact with each other, just as ice and water can at o°. If 
the mixture of the two forms is heated above 95 , all the rhombic 
sulfur becomes monoclinic sulfur. On the other hand, if the 
mixture is cooled below 95 all the monoclinic sulfur changes to 

375 



376 Introduction to General Chemistry 

rhombic sulfur. If the monoclinic sulfur is cooled quickly to 
room temperature the change to rhombic sulfur goes on very 
slowly, just as the change of white to red phosphorus goes on 
slowly at ordinary temperatures, though it is rapid at higher 
temperatures. It is not possible to heat a solid substance above 
its melting-point without melting it, but it is possible to heat a 
solid above its transition point to another solid form and to 
determine the melting-point of the first form, provided the opera- 
tion is carried on very rapidly. In this way it has been found 
that rhombic sulfur melts at 114 . Monoclinic sulfur melts 
at 119 . 

Still another form of sulfur, plastic sulfur, is prepared by 
heating sulfur to its boiling-point and then pouring it into cold 
water. Sulfur so treated is a plastic solid, dark amber in color, 
and when free from rhombic sulfur, which may be formed with 
it, it is insoluble even in carbon disulfide. In time it changes to 
rhombic sulfur. 

With rising temperature molten sulfur is at first a mobile, 
straw-colored liquid, but at about 160 , curiously enough, it 
becomes dark brown and extraordinarily viscous. But when its 
temperature has reached 260 , its color becomes a yellow red, 
and its viscosity again decreases. These changes are due to the 
formation of alio tropic liquid forms. This is the first case of the 
kind which we have encountered. 

601. Chemical Properties of Sulfur. — Sulfur unites chemi- 
cally with most metals and non-metals. Thus it not only forms 
sulfides of metals, but it also forms oxides, chlorides, bromides, 
phosphides, etc. For example, when it is rubbed together with 
mercury, black mercuric sulfide forms : 

Hg+S->HgS. 

If it is heated with iron powder, the reaction to form ferrous 
sulfide begins and soon gives out enough heat to raise the whole 
mass to incandescence (339). The oxides of sulfur are of such 
importance that their chemistry will be described later in detail. 
Chlorine passed over melted sulfur forms sulfur monochloride, 



Sulfur and Its Compounds 377 

S 2 C1 2 . The latter substance, which is a liquid at ordinary 
temperatures, is an important solvent for sulfur. It has been 
produced in enormous quantities for the preparation of the 
terrible mustard gas (695). 

602. The Sources of Sulfur. — Sulfur is found in large deposits, 
supposedly of volcanic origin, in Sicily, Japan, and elsewhere. 
Of these deposits, those of Sicily supplied most of the world's 
trade up to 1904. At that time two mines in Louisiana and one 
in Texas became of tremendous importance, owing to the develop- 
ment of the Frasch process of getting out the sulfur. The 
deposits now being mined there are several hundred feet below 
the surface. Iron pipes are driven down into the beds, and 
superheated steam is forced down to melt the sulfur, which 
collects in a pool at the end of the pipe. From this pool it is 
forced to the surface by compressed air. Single wells deliver 
per day as much as four hundred to five hundred tons of sulfur, 
99 . 5 per cent pure, and keep it up for months at a time. The 
sulfur issuing at the surface is piped into bins, where it is spread 
out in thin layers, which immediately solidify. In this way great 
blocks of sulfur are built, sometimes as much as 150 feet wide, 
250 feet long, and 60 feet high. The greater part of the sulfur 
carried to the coast from the mines is transported in dump- 
bottom cars, much as coal is transported. 

A very considerable quantity of sulfur is produced as a 
by-product of other industries. 

603. Rock or Roll Sulfur, Flowers of Sulfur. — Sulfur from 
the mines of industry is called roll or rock sulfur. Sometimes 
the vapor of sulfur is quickly cooled to produce a powder called 
flowers of sulfur. It is usually contaminated with traces of 
sulfurous and sulfuric acid. 

604. Commercial Importance of Sulfur. — One of the very 
strikingly important uses of sulfur is found in the process of 
vulcanizing rubber. Experts tell us that without sulfur there 
would be no rubber industry worthy of the name. It seems that 
crude rubber is stiff and hard at the temperature of ordinary 
winter weather, and at moderate summer heat it becomes 
very sticky. But after vulcanizing, a treatment with sulfur, 



378 Introduction to General Chemistry 

commonly as vapor or with antimony sulfide (red rubber), it 
gains greatly in strength and elasticity and becomes remarkably 
indifferent to heat. 

Free sulfur is an important constituent of many germicide 
and of many fungicide mixtures. Finely divided sulfur is 
dusted on vines and hops. Before the war Europe used more 
than one hundred thousand tons per annum for this purpose 
alone. Some sulfur is used in making gunpowder, the so-called 
"black powder." Enormous quantities of sulfur are used to 
make sulfur compounds such as sulfur dioxide, sulfites, sulfuric 
acid, carbon disulfide, and a host of other important sub- 
stances. Before the war the annual consumption of sulfur in 
the United States was about three hundred thousand tons, but 
during the war it was reported to be as high as nine hundred 
thousand tons. 

605. Hydrogen Sulfide. — The important reagent hydrogen 
sulfide will have been used in the laboratory by the time the 
work has been carried to this point. A brief description of its 
physical properties and chemistry has already been given (339). 
Usually generators of this gas are available in every laboratory. 
Many special types have been designed, but in principle they 
are much like the hydrogen generators already described in 
294. Ferrous sulfide and hydrochloric acid are almost uni- 
versally used to charge them. 

Hydrogen sulfide burns in air to form water and sulfur 
dioxide. If, however, the flame is allowed to play on a cool 
surface it deposits sulfur. Either as a gas or dissolved in water 
it is a very powerful poison, small doses causing insensibility or 
even death. Fortunately its very bad odor warns of its presence, 
so that it can usually be avoided. 

606. Water Solution of Hydrogen Sulfide. Hydrosulfurous 
Acid. — One volume of water at o° dissolves eighteen volumes of 
the gas, but only about three volumes at room temperature. At 
the latter temperature and atmospheric pressure the saturated 
solution is approximately one-tenth molar in concentration. 
All the hydrogen sulfide may be removed from a solution by 
boiling it. 



Sulfur and Its Compounds 379 

The solution will turn neutral litmus faintly pink. Hydrogen 
sulfide, or hydrosulfurous acid, ionizes in stages as follows : 

H 2 S-»H++HS-, 
HS-->H+ + S— . 

The first or primary dissociation is weaker than that of carbonic 
acid or of dihydrogen phosphate ion, H 2 P0 4 ~ (592). The 
secondary dissociation is extremely small. H 2 S reacts with 
sodium hydroxide to form two products, sodium hydrogen sulfide 
and sodium sulfide: 

H 2 S+NaOH->NaHS+H 2 0, 
NaHS+NaOH->Na 2 S+H 2 0. 

607. Sodium and Ammonium Sulfides, Yellow Ammonium 
Sulfide. — The first of these reactions is nearly complete, but the 
second is far from being so. Crystalline sodium sulfide, Na 2 S, 
can, however, be prepared. A water solution of sodium hydrogen 
sulfide, NaHS, is slightly alkaline, but a solution of the other salt 
is strongly alkaline. These are the results which we would 
expect from the weakness of the acids H 2 S and HS~ (436). 
Solutions of sodium sulfide find important use as a depilatory 
for the removal of hair from the skins of slaughtered animals and 
for other similar purposes. 

The ammonium salts, NH 4 HS and (NH 4 ) 2 S, are prepared by 
passing hydrogen sulfide into ammonium hydroxide solution. 
The resulting solution, which is of course alkaline in reaction, 
is extensively used in chemical analysis. 

Although pure ammonium sulfide solutions are colorless, the 
ordinary reagent is usually yellow, owing to the presence of 
dissolved sulfur formed because of exposure to the oxidizing 
influence of the air. Sulfur dissolves in ammonium sulfide 
forming a series of complex ions which, because of their similarity 
to peroxide ion, are called persulfide ions. Ions of the formula 
S 2 ~ _ , S 3 ", etc., are known. The laboratory reagent, yellow 
ammonium sulfide, which is prepared by dissolving sulfur in 
ammonium sulfide, probably contains a mixture of persulfides. 
Its formula is usually given (NH 4 ) 2 S X , the subscript % denoting 



380 Introduction to General Chemistry 

an undetermined proportion of sulfur. Upon acidification of 
this reagent, sulfur is precipitated and hydrogen sulfide given off: 

(NH 4 ) 2 S,+ 2 HCl->2NH 4 Cl+H 2 S a; , 
H a S*-»H 2 S+(z-i)S. 

608. The Precipitation of Sulfides. — The precipitation of 
sulfides from solution is a very important procedure in making 
an analysis of a substance of unknown composition. The pro- 
cess of a general metal analysis depends on the separation of 
metals from the mixture in order that they may be recognized 
by their characteristic reactions. Thus if we have in solution 
copper sulfate and zinc sulfate, we may separate the metals as 
follows: First, the solution is made acid (0.25 to 0.5 N) with 
hydrochloric acid, and hydrogen sulfide is passed in. Copper 
sulfide alone is precipitated, because the excess of acid makes 
the concentration of sulfur ion so small (432) that only the 
extremely insoluble copper sulfide forms in sufficient amount to 
exceed its molecular solubility (445). Next, the filtrate, which is 
practically free from copper, is made alkaline with ammonium 
hydroxide. Immediately the separation of the white zinc sulfide, 
from solution begins and is carried to completion by the further 
addition of ammonium sulfide. In practice the extremely 
insoluble sulfides of mercury, lead, bismuth, cadmium, arsenic, 
antimony, tin, etc., if present, are brought down with the 
copper sulfide, and then further separation of these metals from 
each other follows. With zinc sulfide appear the sulfides of iron, 
cobalt, nickel, and manganese, together with certain insoluble 
hydroxides. The separation by means of hydrogen sulfide is 
therefore one of the fundamental processes in the analysis of 
metals. 

609. Aqueous Hydrogen Sulfide as a Reducing Agent- 
Aqueous hydrogen sulfide is a good reducing agent by virtue of 
the easily oxidizable sulfur ion, S~~, which it furnishes (503). 
Thus when it is oxidized by the air (339) the fundamental 
change may be represented as follows: 

2S--+0 2 ->2S+20"-. 



Sulfur and Its Compounds 381 

The hydrogen ions which had been associated with the sulfur ion 
before the change unite with the oxygen ion produced, forming 
water. 

Since hydrogen sulfide is acted upon in water solution by 
such mild oxidizing agents as the oxygen of the air, it is not 
surprising to find that it is attacked by very powerful agents 
such as acid permanganate (343) or acid dichromate solutions 
(346). If dilute sulfuric acid is added to potassium permanga- 
nate solution, and hydrogen sulfide is passed into the mixture, 
the intense color of the permanganate soon disappears, and in 
its place we have the nearly colorless solution of manganous 
sulfate. The sulfur ion may be oxidized to free sulfur or to 
sulfite ion or sulfate ion according to the experimental conditions. 
If free sulfur is the main product, as shown by a fine, almost 
white suspension in solution, the fundamental change (561) in 
charges of the atoms undergoing oxidation and reduction is as 
follows : 

2Mn?++5S--->2Mn+++5S, 

and the final equation is 

2 KMn0 4 + 5H 2 S+ 3 H 2 S0 4 -> 2MnS0 4 +K 2 S0 4 + 5S+ 8H 2 0. 

When hydrogen sulfide is led into potassium dichromate solu- 
tion mixed with a strong acid, sulfuric acid for instance, the 
characteristic color change from orange to green or violet of the 
latter reagent on reduction is soon observed. In this case also 
free sulfur, sulfurous acid, or sulfuric acid may be produced by 
the reaction. If the fine, almost white, precipitate of sulfur is 
observed the fundamental equation is 

2Cr 6 ++ 3 S--->2Cr++++ 3 S, 

and the final equation is 

K 2 Cr 2 7 + 4 H 2 S0 4 -f-3H 2 S->K 2 S0 4 +Cr 2 (S0 4 ) 3 +3S+7H 2 0. 

610. Sulfur Dioxide. — Sulfur dioxide is a sharp-odored, 
colorless gas formed when sulfur or sulfides are burned in air. 
It is often found in volcanic gases and is to be noticed in the air 



382 Introduction to General Chemistry 

about towns, since much of the coal burned contains traces of 
sulfur. A 99 per cent pure gas is made on a huge scale by burning 
sulfur, but an enormous supply is also obtained by the burning 
of pyrite, FeS 2 , one of the richest of the sulfur ores: 

4FeS 2 + 1 i0 2 -> 2Fe 2 3 +8S0 2 . 

Since the gas is very easily liquefied it is generally marketed in 
that form and stored in iron tanks, the material of which it 
does not attack. At ordinary temperatures the pressure in 
these tanks is about 2 . 5 atmospheres. The gas dissolves easily 
in water, about 70 volumes to one of water at o° and about 40 
volumes at room temperature. A saturated solution at 20 
and one atmosphere pressure is about a 1 . 6 molar solution. 
The sulfurous acid of commerce contains about 6 per cent of 
sulfur dioxide. 

611. Water Solution of Sulfur Dioxide, Sulfurous Acid. — In 
water solution sulfur dioxide forms sulfurous acid (340) just as 
carbon dioxide forms carbonic acid (152). Sulfurous acid is a 
moderately weak dibasic acid. Its reducing power has already 
been described (340) . On this account it is used on an enormous 
scale commercially as "antichlor," that is, to remove the residual 
hypochlorous acid left after the bleaching action of chlorine. 
Sulfurous acid is also used to bleach substances which chlorine 
rots, such as wool, silk, etc. Sulfurous acid and sulfites are used 
in breaking up the fibers of wood pulp preliminary to paper- 
making. Sulfur dioxide is also used in the preservation of food, 
particularly dried fruits. 

612. Salts of Sulfurous Acid, the Sulfites. — Sulfurous acid 
forms both acid and normal salts. Most of the acid salts are 
soluble, but of the normal salts only the sodium, potassium, and 
ammonium sulfites are soluble in water. Like carbonates (461) 
sulfites interact with strong acids to form the correspondingly 
strong acid salt and the weak and unstable acid, which decom- 
poses, giving off the acid anhydride: 

BaS0 3 +2HCl-»BaCl 2 +H 2 S0 3 , 
H 2 S0 3 ->H 2 0+S0 2 . 



Sulfur and Its Compounds 383 

This, of course, means that sulfites, which are only moderately 
insoluble, are dissolved extensively by acids, particularly if 
the sulfur dioxide is removed from solution by boiling (463). 

613. Sulfur Trioxide. Fuming Sulfuric Acid. — Sulfur tri- 
oxide, S0 3 , the anhydride of sulfuric acid, is a choking gas. It 
dissociates into sulfur dioxide and oxygen more and more with 
increasing temperature; at 400 this dissociation is only about 
2 per cent of the whole; at 500 it is about 9 per cent; at 1000 
it is virtually complete. Sulfur trioxide can be condensed into 
needle-shaped crystals which melt at 18. 8° and form a liquid 

. which may be distilled undecomposed at 44 . 8°. In the presence 
of a trace of moisture which, of course, forms sulfuric acid, sulfur 
trioxide changes to a white, asbestos-like mass which can be 
melted at 50 . This substance probably consists^ of molecules, 
each one of which contains several sulfur trioxide molecules. 
Substances the molecules of which are made up of small mole- 
cules of a simpler substance are called polymers. Sulfur trioxide 
is very hygroscopic. When thrown into water it reacts violently 
to form sulfuric acid. 

Sulfur trioxide dissolves in sulfuric acid to form a product 
called fuming sulfuric acid or oleum. 

614. Sulfuric Acid. — Sulfuric acid has a clear title to being 
the most important chemical used in industry. Very few 
chemical industries exist which do not employ this acid in some 
process. During normal times the United States produced 
about three and a half million tons annually, but during the 
Great War the yearly output is reported to have exceeded 
six million tons (see 618). This statement does not include 
enormous amounts made and used without being put on the 
market. 

615. The Manufacture of Sulfuric Acid. — The problems of 
the manufacture of sulfuric acid are much like those of the 
synthesis of ammonia (520). At low temperatures the union of 
sulfur dioxide with oxygen to form sulfur trioxide, the anhydride 
of sulfuric acid, will go to completion, but too slowly for the 
process to be profitable. If higher temperatures are used the 
rate of the reaction is increased, but equilibrium is established 



384 Introduction to General Chemistry 

with less of sulfur trioxide formed the higher the temperature 
used. The manufacture of the acid has been made successful 
by the use of catalysts, which hasten the reaction at low tempera- 
tures. Two general methods are used. The one called the 
chamber process relies on a gas catalyst and allows the reaction 
to come to completion in huge reaction chambers of lead. The 
other, called the contact process, employs a solid catalyst, on 
the surface of which the reaction proceeds. 

616. The Chamber Process. — In the first method, which is 
the older, the catalyst is a mixture of nitric oxide and nitrogen 
tetroxide. Sulfur dioxide, supplied usually from the burning 
of sulfur or pyrite, is mixed with the right amount of air and is 
then swept over pots which contain niter and sulfuric acid. 
After receiving the nitric acid fumes from the latter the mixed 
gases are sent to the so-called Glover tower to receive more of 
the catalyst, which has been recovered from previous charges in 
a manner which will be described. The reaction begins in the 
Glover tower, but from this tower the gases are forced into the 
great lead chambers in which the reaction is completed. Jets 
of cold water or dilute sulfuric acid in the chambers dissolve the 
sulfur trioxide, forming a moderately concentrated sulfuric acid, 
called chamber acid: 

S0 3 +H 2 0->H 2 S0 4 . 

The actual mechanism by which the catalyst acts is not com- 
pletely understood. The mixture of NO and N0 2 undoubtedly 
contains a little N 2 3 (554) in equilibrium with the former. If 
nitrous acid anhydride is regarded as the active agent, the reac- 
tion may be represented by the following equation: 

N 2 3 +S0 2 ->S0 3 +2NO. 

The nitric oxide left combines with more oxygen, and the regen- 
erated oxides are again able to oxidize more sulfur dioxide. 

The gases left in the chambers are next forced through the 
Gay Lussac tower. The latter is filled with broken brick, over 
which concentrated sulfuric acid constantly trickles. In this 
both sulfur trioxide and nitric oxide dissolve, the first to form 



Sulfur and Its Compounds 385 

fuming sulfuric acid (613) and the second to form nitrosyl 
sulfuric acid: 

4NO+0 2 +4H 2 S0 4 ->4NO-HS0 4 +2H 2 0. 

The mixture is drawn off and carried to the top of the Glover 
tower. From this point it is allowed to run down over the brick 
with which the tower is filled. As it does so it meets a stream 
of dilute acid. The dilution of the nitrosyl sulfuric acid effects 
its decomposition, and oxides of nitrogen are given off to the 
gases entering the tower from the niter pots : 

2NO • HS0 4 +H 2 -> 2H 2 S0 4 +N0 2 +NO. 

The nitric oxide is of course changed into higher oxides by the 
oxygen of the air, so that the catalyst is being continually 
re-formed. The mixed gases then go on the next trip through 
the lead chambers. A little fresh catalyst is furnished con- 
tinually from the niter pots to make up a small unavoidable loss 
in the process. The object of the broken brick in the two towers 
is to afford a large surface of contact for the gases. The yield 
of sulfuric acid is about 95 per cent of the amount theoretically 
possible. The chamber acid is of a concentration of about 
78 per cent H 2 S0 4 and 22 per cent water. More concentrated 
acid is made by subsequent evaporation. 

The following diagram, a so-called flow sheet, shows by 
means of arrows how the various materials entering into and 
formed in the chamber process are related to one another : 

Sulfur Air 



1 

so 2 


N 2 3 




1 

H 2 S0 3 


NO 


Air 


1 
Dil. H 2 SO 


4 N 2 3 


Cone. H 2 S0 4 


1 \ 
Cone. H 2 S0 4 Dil. 


H 2 S0 4 


1 
NO-HS0 4 




H 2 S0 4 


N 2 3 



3 86 



Introduction to General Chemistry 



Sulfur, air, water, and the oxides of nitrogen are the start- 
ing materials. Of these the oxides of nitrogen, which act 
catalytically, are largely recovered in the process and used over 
and over again. The small deficiency caused by the failure to 
recover them completely is made up by gases from concentrated 
nitric acid. 

617. The Contact Process. — The more recent method of 
making sulfuric acid is the so-called contact process. Sulfur 
dioxide is made to unite with oxygen on the surface of a catalyst 
which can be iron oxide, nickel oxide, brick, quartz, etc. 
Platinum is an especially good agent for this purpose, since it 
allows the reaction to be carried on at temperatures as low as 
450-500 . Other catalysts are reported to be in use. If platinum 
is employed the gases must be carefully freed from arsenic, since 

TABLE XXV 



Be 

(Degrees Baume) 


(DegreelTwadell) S P ecific Qr ^^ 


Percentage of 
Sulfuric Acid' 


5o° 

60 

66 


I05° 

141 

167 


i-53 
1. 71 
1.84 


62 
78 
93 



this substance, which is commonly present in pyrite, renders 
the platinum ineffective. The sulfur trioxide formed is dis- 
solved in concentrated sulfuric acid, and water is added to the 
product to make the more dilute acids. The contact method has 
the advantage that much more concentrated solutions of sulfuric 
acid may be prepared directly by it than by the chamber process; 
but in spite of this advantage and the greater simplicity of the 
method, the chamber process is still the more commonly used, 
owing to many skilful improvements in the details of manu- 
facture. 

618. The Baume and Twadell Scales of Specific Gravity. — 
In the market reports, etc., the concentration of sulfuric acid is 
usually given in terms of either of two arbitrary scales of specific 
gravity: in England and the United States the Baume (abbre- 
viation Be) and on the Continent the Twadell (abbreviation Tw). 
Table XXV gives some data for the temperature of 15 . 



Sulfur and Its Compounds 387 

• The " concentrated sulfuric acid" of the laboratory is usually 
understood to be 66° Be. The annual production of sulfuric acid 
is given in terms of 66° acid (614). The so-called chamber acid 
(616} is usually 6o° Be. The specific gravity of other liquids 
heavier than water is also given on these scales. Another 
Baume scale is used for liquids lighter than water. Complete 
tables showing the relationship between the specific-gravity 
scales will be found in most collections of chemical tables. 

619. Physical Properties of Sulfuric Acid. — Pure sulfuric acid 
is a heavy, oily liquid often called oil of vitriol. Like hydro- 
chloric acid it forms with water a mixture of constant boiling- 
point (251). At one atmosphere pressure the composition of 
this mixture is 98.4-6 per cent acid. It boils at 3 3 8°. Pure 
sulfuric acid boils at 290 . The vapor of sulfuric acid consists 
almost entirely of sulfur trioxide and water. 

Sulfuric acid is very hygroscopic and therefore a good drying 
agent. Since the 98 . 4 per cent acid does not give off fumes at 
ordinary temperatures, a little is often kept in storage vessels, 
such as desiccators, to keep the air dry. If a moist substance is 
put in such a vessel it will continuously lose water to the air, 
and as the water vapor is constantly being taken up by the 
sulfuric acid the substance is soon dry. The great heat which is 
given out as sulfuric acid is diluted with water has already been 
mentioned. Students have been warned (93) that there is 
danger from spattering whenever sulfuric acid is mixed with 
water. On this account the acid is always poured very slowly 
into water, with gentle stirring in order that the heavy layer may 
be well mixed into the lighter layer. The process of pouring 
water into sulfuric acid is very likely to result in a violent 
spattering of the hot and corrosive acid. 

620. Chemical Properties of Sulfuric Acid. — Dilute sulfuric 
acid is a dibasic acid a little weaker than the strongest acids : 

H 2 S0 4 ±5H++HS0 4 -, 
HS0 4 -^H++S0 4 --. 

Even its secondary dissociation is that of a moderately strong 
acid. Like other acids, dilute sulfuric acid reacts with metals 



388 Introduction to General Chemistry 

above hydrogen in the electromotive series (492) to give off 
hydrogen, but the sulfate ion is usually very slowly affected by 
reducing agents. 

On account of its high boiling-point concentrated sulfuric 
acid, contrary to the usual rule, can be made to displace stronger 
acids from their salts, since the acid formed by the partial double 
decomposition of sulfuric acid with the salt is distilled off at 
temperatures at which the sulfuric acid is but little volatile. 
In this way even the insoluble silver chloride may be dissolved, 
though the action is very slow : 

AgCl(solid)^AgCl(dissolved)^Ag++Cl- 

* H 2 S0 4 ^HS0 4 -+H+ 
IT It 

AgHS0 4 HC1 

Concentrated sulfuric acid, unlike the dilute, is an oxidizing 
agent as well as an acid, though its power to oxidize is not as 
strong as is that of nitric acid. Some of the important and 
characteristic reactions will be discussed shortly (622). 

If concentrated sulfuric acid is put on filter paper the latter 
will be found to blacken. This is a typical behavior of sulfuric 
acid with any carbon compound from which it can take water 
from combination. Thus sugar, C I2 H 22 Ou, is charred by sulfuric 
acid. For the same reason sulfuric acid burns the skin. These 
reactions are in reality more complex than simple dehydrations. 

621. Concentrated Sulfuric Acid and Metals. — Sulfuric acid 
does not attack silver or gold, but the impure acid does attack 
platinum. In the cold it does not react with the less active 
metals like copper, but does do so when hot, forming the sulfate 
of the metal and sulfur dioxide. The more active metals like 
zinc are attacked by the warm acid with the liberation of S0 2 ; 
but sulfur and hydrogen sulfide may be formed according to 
experimental conditions. That sulfuric acid in the absence of 
air does not attack iron is of great technical importance, since 
the latter material is widely used for containers of this acid. 

622. Oxidation by Concentrated Sulfuric Acid. — The solu- 
tion of metals in sulfuric acid is of course the result of the oxidizing 



Sulfur and Its Compounds 389 

action of the acid. Table XXVI states briefly the charges 
which the sulfur atom carries in sulfuric acid and its reduction 
products, as well as in the other common compounds which we 
have studied. Apparently a sulfur atom in sulfuric acid must 
gain two electrons to become sulfite, six to become free sulfur, 
and eight to become hydrogen sulfide. 

TABLE XXVI 

Compounds Sulfur Atoms 

S0 3 , H 2 S0 4 , and sulfates S 6 + 

■ S0 2 , H 2 S0 3 , and sulfites S<+ 

Free sulfur S° 

H 2 S ~ . . . . . S — 

Of the possible reduction products of sulfuric acid, S and 
H 2 S are capable of further reaction with excess acid. Thus 
H 2 S bubbled through concentrated H 2 S0 4 forms sulfur and the 
unstable sulfurous acid: 



H 2 S+H 2 S0 4 ->S-r-H 2 0+H 2 SO 



Sulfur warmed with concentrated H 2 S0 4 reduces the acid, form- 
ing S0 2 : 

S+2H 2 S0 4 ->S0 2 +2H 2 S0 3 . 

As in the case of nitric acid (557), so with sulfuric acid, the 
concentration of the acid, the temperature of the reaction, the 
fineness of division of the reducing agent if it is a solid, and 
whether or not the acid is present in excess, all affect the nature 
of the products. 

'In the case of zinc the fundamental reactions by which 
each of the possible products is formed are as follows: 

for S0 2 , Zn+S 6 + ->Zn++ + S4+, 
for S, 3Zn+S 6 + -» 3 Zn+++S°, 
forH 2 S, 4 Zn+S 6 +->4Zn++ + S--, 

and the final equations are 

Zn+ 2 H 2 S0 4 -> ZnS0 4 + S0 2 + 2 H 2 0, 
3Zn+4H 2 S0 4 ->3ZnS0 4 +S+4H 2 0, 
4 Zn+5H 2 S0 4 -> 4 ZnS0 4 +H 2 S+4H 2 0. 



390 Introduction to General Chemistry 

623. Action of Non-Metallic Reducing Agents on Concen- 
trated Sulfuric Acid. — The non-metallic reducing agents show 
the same differences in behavior with concentrated sulfuric acid 
as do the metals. Thus hydrogen bromide reduces sulfuric acid 
to form sulfurous acid and bromine, while hydriodic acid reduces 
sulfuric acid to sulfur and hydrogen sulfide, iodine and water 
being formed at the same time (341). The action of sulfur and 
hydrogen sulfide respectively has just been discussed (622). 
Charcoal (carbon) reduces the acid to sulfurous acid and is itself 
oxidized to carbon dioxide: 

C+ 2 H 2 S0 4 -> C0 2 + 2 H 2 S0 3 . 

624. Sulfates. — Solutions of the acid sulfates are acid to 
litmus, owing to the fact that the HS0 4 ~ ions are very largely 
dissociated in solutions of even moderate dilution. Sulfates 
as a rule are easily soluble in water, but those of barium, stron- 
tium, calcium, and lead are little soluble. Under dry heat the 
acid sulfates of sodium and potassium first lose water, forming 
the so-called pyrosulfates. These on further heating change 
into the normal sulfates with the loss of sulfur trioxide : 

2Na(HS0 4 ) 2 ->Na 2 S 2 7 +H 2 0, 
Na 2 S 2 7 ^ Na 2 S0 4 + S0 3 . 

Normal sulfates of the alkalies are very resistant to heat. The 
others break down, forming the corresponding oxides and sulfur 

trioxide : 

Fe 2 (S0 4 ) 3 ->Fe 2 3 + 3 S0 3 . 

625. Sodium Thiosulfate. — One other sulfur compound 
should be mentioned. This is sodium thiosulfate, the "hypo" 
used in photography to " fix" negatives. If sulfur is boiled in a 
solution of sodium sulfite it will soon be seen to dissolve. Upon 
evaporation the solution gives clear, colorless crystals of sodium 

thiosulfate : 

Na 2 S0 3 +S^Na 2 S 2 3 . 

This salt is extremely soluble in water. At o°, 217 parts dissolve 
in 100 parts by weight of water. It will be remembered as one 



Sulfur and Its Compounds 391 

of the substances used to show the phenomenon of supersatura- 
tion of a solution (123). When solutions of hypo are made, acid 
sulfur is deposited and sulfurous acid forms: 

Na 2 S 2 3 + 2HCI -> 2 NaCl+H 2 S0 3 +S. 

Sodium thiosulfate is manufactured in immense quantities for 
purposes similar to those for which sodium sulfite is used. The 
reaction with chlorine is as follows : 

Na 2 S 2 3 + 4 Cl 2 + 5H 2 -> 2 NaCl+ 2 H 2 S0 4 + 6HC1. 

Sodium thiosulfate reacts differently with iodine, forming sodium 
tetrathionate and sodium iodide : 

2Na 2 S 2 3 +I 2 -> 2NaI+Na 2 S 4 6 . 

626. The Silver Thiosulfate Complex Ion. — If a solution of 
sodium thiosulfate is added to a solution of silver nitrate a white 
precipitate of silver thiosulfate appears: 

2 AgN0 3 +Na 2 S 2 3 -> Ag 2 S 2 3 + 2NaN0 3 . 

The precipitate dissolves very easily in excess of the hypo solu- 
tion, forming complex silver thiosulfate ions, Ag 2 S 4 06~~, and 
secondarily some of the sodium silver thiosulfate molecules : 

Ag 2 S 2 3 +Na 2 S 2 3 ->Na 2 Ag 2 S 4 6 . 

If we add sodium chloride to the resulting solution no precipitate 
of silver chloride will appear. Or if we treat silver chloride with 
sodium thiosulfate solution the former will dissolve. So little 
dissociated is the complex silver thiosulfate ion that even silver 
bromide, which is less soluble than silver chloride, is dissolved 
by the hypo solution. 

627. The Action of Hypo as a Fixing Solution. — The coating 
of a photographic plate or film which is sensitive to light is 
ordinarily made of gelatine and silver bromide. After the plate 
has been "exposed" and "developed" the parts acted upon by 
light are black coatings of finely divided silver. The unchanged 
part of the "negative" is still coated with the light-yellow silver 
bromide, which must be removed before the former is again 



392 Introduction to General Chemistry 

exposed to light, or the whole plate will become evenly black 
and the picture be obliterated. Accordingly the unchanged silver 
bromide is dissolved in the hypo bath, and the soluble sodium 
silver thiosulfate thus formed is washed away. 

628. Other Sulfur Compounds. — The chemistry of sulfur and 
its compounds has not been exhausted by any means by the dis- 
cussions of this chapter. The treatment here has been limited 
to the more important properties of the commonest sulfur com- 
pounds; but there are numerous others, some of which are of 
technical importance. Of the latter class are persulfuric acid, 
H 2 S 2 08 permonosulfuric or Caro's acid, H 2 S0 5 ; and hyposulfur- 
ous acid, H 2 S 2 4 , together with their salts. The first two of 
these acids are closely related to hydrogen peroxide. Their 
graphic formulae (323), together with that of sulfuric acid, are 
thought to be as follows: 

HO-S0 2 HO-S0 2 HO-S0 2 

I • I ! 

000 

I i * ! 

O O H 

i I 

HO-S0 2 H 

Persulfuric acid Permonosulfuric acid Sulfuric acid 

They and their salts are powerful oxidizing agents. The Zn 
salt of hyposulfurous acid is readily formed by dissolving zinc 
dust in sulfurous acid: 

Zn+ 2ILSO3 ~> ZnS 2 4 + 2H 2 0. 

The sodium salt Na 2 S 2 4 , known in commerce as sodium 
hydrosulfite, is a powerful reducing agent and has strong bleach- 
ing action on vegetable colors. It is used for bleaching soap, etc. 
Five other acids of theoretical but at present little practical 
importance are the following: dithionic acid, H 2 S 2 6 ; trithionic 
acid, H 2 S 3 06; tetrathionic acid, H 2 S 4 6 ; pentathionic acid, 
H 2 S 5 6 ; hexathionic acid, H 2 S 6 6 . 



CHAPTER XXV 
CARBON AND CARBON COMPOUNDS: ORGANIC COMPOUNDS. I 

629. Review. — Some relatively simple reactions of carbon 
and carbon compounds have already been studied. When 
carbon is burned in oxygen, the gas carbon dioxide is formed 
and may be identified by its reaction with limewater to form a 
white precipitate of calcium carbonate (18, 151). The burning 
of other carbon compounds in excess of air is found to produce 
the same gas (20). Charcoal heated with copper oxide reduces 
the latter, forming free copper and carbon dioxide (328). The 
commercial production of phosphorus depends on the reducing 
action of carbon on calcium phosphate mixed with sand (580). 
Carbon monoxide is also a good reducing agent at high tempera- 
tures (329). When carbon dioxide dissolves in water the weak 
and unstable carbonic acid is formed (152, 285, 449). The 
bicarbonate and carbonate of sodium and of potassium are 
common reagents (161, 162, 448). Limestone and marble, 
different forms of calcium carbonate, are used on a huge scale 
for the preparation of both lime (150) and carbon dioxide (163). 
The precipitation of carbonates is governed by the usual condi- 
tions which control the precipitation of little soluble salts of 
weak acids (448). All carbonates are soluble in strong acids 
(461). Another common reagent, acetic acid, is also a carbon 
compound (157). It is a colorless liquid, of sour odor, miscible 
with water in all proportions. It is used mainly as a weak and 
stable acid (424). 

630. Physical Properties of Carbon. — The free element car- 
bon appears in several allotropic forms. An impure form of 
great industrial importance is the non-volatile portion of coal 
left after the latter has been subjected to processes which free 
it of volatile matter. This material is called coke. It is 
employed as a fuel directly but is also used for making different 
forms of fuel gas, which are usually mixtures of carbon monoxide 

393 



394 Introduction to General Chemistry 

with other gases. Coke is also used for the reduction of oxide 
ores such as the important iron ore hematite, Fe 2 3 . Thus 
coke is a necessary material for many industries. 

Graphite is another form of carbon. It occurs to some extent 
in nature but is now produced in enormous quantities by heating 
coke in the electric furnace (Acheson process). It is formed 
in slippery black scales which serve as an excellent lubricant 
for parts exposed to high temperatures and for wood surfaces 
which must rub together. Since graphite conducts electricity 
it is an important material for electrodes, particularly for the 
preparation of substances like chlorine. Mixed with clay, 
graphite is the "lead" of lead pencils. 

The diamond is pure crystalline carbon. The natural 
crystals are cut into many-faced shapes in order to bring out 
the power of this material to reflect light. Diamonds were 
first produced artificially by Moissan. He took advantage of 
the fact that carbon, though insoluble in all ordinary solvents, 
is soluble in molten iron, and dissolved graphite in the latter 
substance. He then chilled the mass quickly by plunging it 
into molten lead. The outside layer of iron was first to cool, 
and as it shrank in so doing it put the inner part under great 
pressure. When the entire mass was cooled tiny diamonds, 
about 0.5 mm. in diameter at the most, were found in the 
interior. The diamond has the distinction of being the hardest 
substance known and is consequently an important abrasive 
and cutting material. Black diamonds, which have no orna- 
mental value, are set in the cutting surfaces of rock drills, while 
small diamond chips are used to cut glass. 

631. Chemical Properties of Carbon. — Any form of carbon 
when burned gives either carbon monoxide or carbon dioxide, 
according to the supply of oxygen available. 

If sulfur vapor is led over hot carbon, carbon disulfide, CS 2 , 
is formed." Carbon disulfide is a colorless liquid. It • is an 
important solvent for a variety of substances. An instance of 
interest is its use to dissolve rubber. As has already been 
pointed out, it is very easily inflammable and therefore danger- 
ous (546). 



Carbon and Carbon Compounds 395 

Carbon does not unite with hydrogen (except to a very small 
extent at high temperatures) without the assistance of a catalyst; 
but if hydrogen is led over a mixture of carbon and very finely 
divided nickel at 250 a very good yield of the gas methane, 
CH 4 , is obtained. 

Carbon can be made to unite with metals, forming carbides. 
Thus when coke is heated with lime in the electric furnace 
calcium carbide is formed: 

CaO+ 3 C->CaC 2 +CO. 

This substance is of interest because of the part it plays in the 
cyanamide process for fixing atmospheric nitrogen (526). It 
is also of importance because of its ready action with water to 
form acetylene: 

CaC 2 + 2H 2 -> Ca(OH) 2 + C 2 H 2 . (49) 

When sand is heated with coke under special conditions car- 
borundum, SiC, is formed. This important abrasive is ground, 
mixed with a binder, and molded into grinding wheels, knife 
sharpeners, etc. 

632. Carbon Monoxide. — Wherever combustion of carbon 

or carbon compounds goes on with a deficient supply of air, the 

odorless, colorless, and very poisonous gas carbon monoxide is 

likely to be formed. It is often seen burning as a pale-blue 

flame on the top of coal fires. Care must be taken that furnaces 

are sufficiently ventilated to prevent this dangerous gas from 

getting into living quarters. Carbon monoxide is made in large 

quantities for industrial fuel gas. If made from air and coke 

it is mixed with nitrogen. Steam and coke give carbon monoxide 

and hydrogen: 

H 2 0+C->CO+H 2 . 

633. Carbon Dioxide. — Liquid carbon dioxide is sold in steel 
cylinders in which it is under about 60 atmospheres' pressure. 
When it is allowed to flow rapidly from these into a cloth bag 
it is cooled so strongly by its own evaporation that tiny particles 
of the solid carbon dioxide are formed and collect in the bag as 
carbon dioxide snow. This substance has a vapor pressure of 



396 Introduction to General Chemistry 

76 cm. at — 79 , a temperature which is 23 below the melting- 
point of the solid. Thus the latter evaporates at — 79 without 
melting. Carbon dioxide snow has a limited but important use 
as a refrigerant. 

The gas is very stable and will not support ordinary com- 
bustion. It is so much heavier than air that it can be poured 
from one vessel to another much as we would pour a liquid. 
These properties make it very serviceable as a fire extinguisher. 
Usually the apparatus contains concentrated carbonate solution 
and a bottle of sulfuric acid. When the handle of the device is 
turned the bottle is broken, and a concentrated solution of 
carbon dioxide is formed under considerable pressure of the gas. 
A stream of the gas and supersaturated solution may then be 
directed against the fire. 

The so-called charged water (soda water) is, of course, a 
solution of carbon dioxide under pressure. When the pressure 
is released the solubility of the gas decreases, and the liquid is 
seen to froth with the escaping bubbles. 

Carbon dioxide plays an important part in plant life, as will 
be shown later (690, 691). 

634. Illuminating Gas. — The distillation of coal yields 
illuminating gas, together with other important products. 
These vary in composition according to the kind of coal used 
and the temperature to which the latter is heated. The follow- 
ing data are representative of the nature and the approximate 
amounts of products gained from one ton of coal: 

Illuminating gas, 11,000 cu. ft. 
Ammonia, 6 lb. 
Coal tar, 120 lb. 
Coke, 1,500 lb. 

The coal is heated in retorts, A, to about 1300 (Fig. 93). 
The vapors given off are sent through water in the so-called 
hydraulic main, B, in which some of the ammonia and tar are 
collected. The gas is then passed first through condensers, 
C, to take out the rest of the coal tar, and next through trays 
of broken brick, D, " scrubbers," over which water is trickling, 



Carbon and Carbon Compounds 



397 



to dissolve the last of the ammonia. Next, objectionable impu- 
rities such as carbon dioxide and hydrogen sulfide are removed. 
Finally the gas is stored in huge holders preliminary to its 
passage into the city main. Illuminating gas consists largely 
of hydrogen and methane, CH 4 , together with a small proportion 
of ethylene, C 2 H 4 (660). 

635. Organic Chemistry. — The number of definite compounds 
containing the element carbon far exceeds that of all other 
chemical substances. It was once thought that all except the 
very simplest carbon compounds were products of vegetable 
and animal organisms, and for this reason they were called 

1 



a^ 






Fig. 93 



organic substances. We now know that this view is erroneous, 
because the great majority of organic substances can at present 
be made by purely chemical methods (synthesized) from the 
elements composing them. Although the term " organic" is 
a misnomer, it is retained and serves to distinguish the great 
class of compounds of carbon from all others, which by contrast 
are known as inorganic substances. A few of the simplest com- 
pounds of carbon are classed as inorganic; these include such 
substances as the oxides of carbon, carbonates, carbides (e.g., 
CaC 2 ), etc. 

It will be the purpose of the rest of this chapter to give the 
student a brief glimpse of the fundamental principles of organic 
chemistry, and the object of the next chapter to show some of 
the great successes of this branch of the science. No attempt 
should be made on the part of the reader to secure an intimate 
knowledge of the detail presented, since such effort is best 



398 Introduction to General Chemistry 

expended under circumstances in which more time is given to 
laboratory work than is possible in a general chemistry course. 
But the student should aim to understand the fundamental 
principles pointed out. 

We shall begin by the study of one of the commonest and 
best known of organic substances, starch. We shall then show 
how from this plant product a variety of other substances can 
be derived by chemical processes. This procedure will lead us 
to a knowledge of the properties and interrelations of several 
of the most important organic compounds and thus pave the 
way for the further study of the compounds of carbon. 

636. Starch. — Starch in nearly pure form is a well-known 
article of merchandise, used extensively as food and also for 

TABLE XXVII 

Percentage of Starch in Several Plant 

Products 

Wheat 68 

Corn (maize) 75 

Oats 55 

Rice 78 

Potatoes 15 

Beans 59 

the stiffening (starching) of laundered clothes, etc. Most starch 
is made from corn (maize), of which it constitutes about 75 per 
cent. Starch forms a large part of the mass of most grains and 
seeds and is present in large amount in tubers, bulbs, and other 
parts of plants. Table XXVII shows the starch contents of a 
variety of plant products. 

Starch may be easily prepared by grating a potato and 
stirring the pulp with water. If the mixture is poured through 
a coffee strainer, the starch goes through with the water, while 
the fiber is left behind. The (insoluble) starch soon settles 
out of the water, which can then be decanted. After being 
dried at room temperature the starch is obtained as a white* 
powder. 

It is well known that starch is one of the indispensable food 
constituents for man and all herbivorous animals. It constitutes 



Carbon and Carbon Compounds 399 

the most valuable ingredient of vegetables and fruits and forms 
a very considerable part of all breadstuff's. 

The simplest formula calculated from the analysis of starch 
is C6H I0 O s ; but since it is not possible to volatilize this sub- 
stance its molecular weight cannot be found by the vapor- 
density method. The physical and chemical properties of 
starch lead us to Relieve that its formula is less simple than 
that given, and that it is better represented by (C 6 H I0 O s ) w , where 
n may be a rather large number. 

637. The Properties of Starch. — Starch occurs in plants in 
minute grains easily visible under the microscope. Fig. 94 
shows the appearance of wheat starch. 

It scarcely need be said that the visible 
grains are not the molecules of starch, for 
each grain contains an enormous number 
of molecules. 

Natural starch is nearly insoluble in 
cold water, but when boiled with water 
the grains burst, producing so-called 
starch paste. A solution of starch does Fig. 94 

not conduct an electric current much 

better than water; it is not an acid, base, or salt. The action 
of iodine on starch (263), giving a blue color, forms the simplest 
and best test for this substance. 

638. The Conversion of Starch into Glucose. — The most 
important reaction of starch is that in which it unites with water 
to form glucose, or grape sugar: 

C 6 H I0 O s +H 2 O->C 6 H 12 O 6 . 

This reaction requires the aid of a catalytic agent. Acids of all 
kinds act as catalytic agents for the conversion of starch into 
glucose. The latter substance is made commercially by heating 
starch with very dilute hydrochloric acid. The acid rapidly 
promotes the union of water with the starch but is not itself 
changed. At the end of the reaction the HC1 is neutralized 
with soda and thus converted into NaCl, the presence of which 
in minute amounts is not objectionable. The union of starch 




400 Introduction to General Chemistry 

with water is called hydrolysis, and we say that starch is hydro- 
lyzed to glucose. 

639. Glucose, C 6 H I2 6 . — The sugar of ripe grapes and of 
many other fruits is largely glucose, which is also called dextrose 
and grape sugar. In pure form glucose is a white crystalline 
substance looking much like ordinary "granulated sugar." It 
is very soluble in water and is about half as sweet as ordinary 
sugar. Very little pure glucose is made, but glucose sirup is 
produced in enormous quantities. 

Glucose sirup is used directly as a table sirup and is com- 
monly known as corn sirup. It is frequently colored and 
flavored to imitate maple sirup. It is a good and cheap food- 
stuff. A large part of the production of glucose sirup is used 
in the manufacture of candy. The sirup is now shipped largely 
in tank cars. 

640. The Fermentation of Glucose and the Formation of 
Alcohol. — The fermentation of fruit juices is due to the decom- 
position of their glucose content into alcohol, C 2 H 6 0, and CO,, 
according to the equation 

C6H I2 6 ->2C 2 H60+2C0 2 . 

This change requires the aid of a catalytic agent produced by 
growing yeast, which is a simple form of vegetable organism 
(Fig. 95). It was once thought that the 
yeast consumed the glucose as food and 
produced alcohol and C0 2 as products. 
The falsity of this idea was shown by 
grinding up growing yeast so as to 
destroy every plant cell, then pressing 
out the plant sap and mixing it with 
glucose. The sap or extract so obtained 
F IGi 95 fermented glucose and therefore con- 

tained the active catalytic agent of fer- 
mentation. Alcoholic beverages such as beer, wine, and whiskey 
are made by the fermentation of liquids containing glucose. 

641. Alcohol, C 2 Ij60. — There are many kinds of alcohols, 
but the term alcohol as commonly and popularly applied to the 




<v 



<# 



Carbon and Carbon Compounds 401 

fermentation product of glucose, C 2 H 6 0. This is also known 
as grain alcohol (because it is usually made from corn, barley, or 
rye as starting materials), and by chemists as ethyl alcohol. 
Alcohol is a colorless liquid which burns with a non-luminous, 
sootless name. It is miscible with water in all proportions. 

In the manufacture of alcohol, corn from which the germ has 
been removed is ground up coarsely and boiled with water. The 
product is cooled to 6o° or 65 and mixed with the malt, which 
furnishes the diastase necessary to convert the starch into 
maltose, Ci 2 H 22 On, which is an intermediate product between 
starch and glucose. Maltose is easily changed into glucose by 
heating with dilute HC1 or H 2 S0 4 : 

C u H M 0„+H 2 -> 2C 6 H I2 6 , 

and like glucose is easily fermented to give alcohol and C0 2 : 

C I2 H 22 II +H 2 0->4C 2 H 6 0+ 4 C0 2 . 

After all starch has changed into maltose, yeast is added to 
start fermentation, which is finished in three days. The product, 
which contains 10 to 12 per cent of alcohol, is then distilled. 
The alcohol (boiling-point 7 8°) distils off more readily than the 
water and is thus easily separated from most of the water as 
well as from the other soluble and insoluble materials present. 
Repeated distillation finally yields a product containing 95 per 
cent of alcohol; this is the ordinary commercial grade, called 
proof spirit. Pure alcohol, free from water, is called absolute 
alcohol. Denatured alcohol is proof spirit to which has been 
added methyl alcohol, benzene, pyridine, etc., to render it unfit 
for drinking. It is poisonous. 

642. Ether, C 4 H I0 O. — Ether is made by allowing alcohol to 
drop slowly into a mixture of alcohol and two parts of concen- 
trated H 2 S0 4 at 140 . Crude ether distils off and is condensed 
by cooling its vapor (Fig. 96). The reaction in its simplest form 
may be represented thus: 

2 C 2 H 6 O->C 4 H I0 O+H,O. 



4-02 



Introduction to General Chemistry 



The H 2 S0 4 acts as a powerful dehydrating (water-absorbing) 
agent. It is very probable that the reaction is not as simple 
as indicated, but that an intermediate compound of C 2 H 6 and 
H 2 S0 4 is first formed and later decomposed. 

Ether is a colorless, mobile, and very volatile liquid boiling 
at 35 . It is our most important anaesthetic, being almost 
universally used in surgical operations. It is also of enormous 
importance chemically. It is an excellent solvent for numerous 
organic substances. It mixes with alcohol in all proportions 
but is only slightly soluble in water. 




Fig. 96 

643. The Paraffine Series. — From crude petroleum and the 

accompanying gases a very remarkable series of carbon and 
hydrogen compounds, called hydrocarbons, can be obtained. 
The simplest member of the series is methane, CH 4 (54). It 
was pointed out that this gas is the chief component of natural 
gas and of marsh gas. Methane may be made in the laboratory 
by heating a mixture of sodium acetate, sodium hydroxide, and 
lime to a high temperature in an iron retort. The reaction is 
substantially as follows: 

NaC 2 H 3 2 +NaOH -> CH 4 +Na 2 C0 3 . 

The second member of the series is ethane, C 2 H 6 , a gas closely 
resembling methane. The third member is propane, C 3 H 8 . 
The formula, name, physical state, boiling-point (B.P.), and 
melting-point (M.P.) of a number of these compounds are 
given in Table XXVIII. It will be seen that the number of 



Carbon and Carbon Compounds 



403 



H atoms is two more than twice the number of carbon atoms in 
any molecule. Therefore the general formula for a compound 
having n carbon atoms is C w H 2W + 2 . All the members of this 
series of compounds up to C 20 are known, and also several with 
more carbon atoms in the molecule. 



TABLE XXVIII 



CH 4 , methane, gas. . . 
C 2 H6, ethane, gas .... 
C 3 Hs, propane, gas. . . 
C4IL.0, butane, gas . . . 
C S H I2 , pentane, liquid 
C6H14, hexane, liquid . 



B.P. 



-164 
" 93° 

- 17° 

+ i° 

38° 

7i° 



C 7 Hi6, heptane, liquid. 
CsHig, octane, liquid. . 

C16H34, liquid 

CzoH^, solid 

C 24 H 5 o, solid 

C32H66, solid 



B.P. 



124^ 



M.P. 



18" 
37° 
Si° 
68° 



The petroleum products in common use are mixtures of 
hydrocarbons. The principal components of some of these are 
given in Table XXIX. 

TABLE XXIX 

Gasoline , C 5 H I2 to C 8 H l8 

Kerosene CgEks to C I2 H 2 6 

Lubricating oils C I2 H 2 6 to Ci6H 34 

Vaseline d6H 34 to C 20 H 42 

Paraffine C 24 Et S o to C 3 6H 74 

644. The Action of Chlorine on Methane. — Chlorine acts on 
methane to form a series of four compounds. The first reaction 
yields methyl chloride, CH 3 C1: 

CH 4 + Cl 2 -> CH3CI+HCI. 

Methyl chloride is a colorless gas which can react still further 

with chlorine: 

CH 3 C1+C1 2 ->CH 2 C1 2 +HC1. 

The product, CH 2 C1 2 , methylene chloride, is a colorless 
liquid, which gives, with more chlorine, chloroform, CHC1 3 : 

CH 2 C1 2 + Cl 2 -> CHCI3+HCI. 

Chloroform is a sweet-smelling, colorless liquid boiling at 6i° 
and insoluble in water, but miscible with alcohol or ether in all 



404 Introduction to General Chemistry 

proportions. It is a good anaesthetic but is not as safe as ether 
for patients with weak hearts. Finally the action of more Cl 2 
on CHCI3 gives carbon tetrachloride: 

CHC1 3 +C1 2 ->CC1 4 +HC1. 

This is a colorless liquid which is not only incombustible but 
an excellent fire extinguisher. It is the principal component 
used in Pyrene and similar fire extinguishers. Devices of this 
type should be at hand in every laboratory for the speedy control 
of fires that might otherwise prove dangerous. 

Carbon tetrachloride and closely related substances are used 
as cleaning fluids. They are perfectly safe, since they are not 
combustible. 

645. Methyl Alcohol, CH 4 0.— If methyl chloride, CH 3 C1, is 

heated with water at a high temperature under pressure it reacts 

as follows: 

CH 3 C1+H 2 -> CH 4 0+HC1. 

The substance, CH 4 0, is methyl alcohol. This method of mak- 
ing methyl alcohol is only of scientific interest. The substance 
is obtained in large quantity as one of the volatile products 
of the distillation of wood (701). The common name of methyl 
alcohol is wood alcohol. It is a colorless liquid, boiling at 66° 
and miscible with water in all proportions It resembles com- 
mon (ethyl) alcohol in its physical and chemical properties. 
It is a dangerous poison, often accidentally causing permanent 
blindness or death. It is a valuable solvent and is manufactured 
on a large scale. 

646. Methyl Ether, C 2 H 6 0. — Methyl ether is made from 
methyl alcohol in practically the same way that ethyl ether, 
C 4 H I0 0, is made from ethyl alcohol: 

2CH 4 0->C 2 H 6 0+H 2 0. 

Methyl ether, QH^O, is a colorless gas having chemical prop- 
erties resembling those of ethyl ether. 

647. Isomerism. — The formula given for methyl ether is 
that based upon analysis and gas density. It will be recalled 






u 



Carbon and Carbon Compounds 405 

that exactly the same formula, C 2 H 6 0, was also ascribed to ethyl 
alcohol. This formula was also fixed by analysis and vapor 
density. Both substances, although entirely distinct in physical 
and chemical properties, have exactly the same percentage com- 
position, the same vapor density, and therefore the same molecular 
weight! This is a most remarkable fact but by no means the 
only known case of the kind. Indeed, among organic com- 
pounds there are hundreds, yes thousands, of cases in which 
two or even several totally different substances have the same 
percentage composition and the same molecular weight and 
in consequence the same formula. Such pairs or groups of 
substances are called isomers. In consequence we speak of 
the isomerism of methyl ether and ethyl alcohol. We shall 
next consider some theoretical matters with the object of 
developing an explanation of the remarkable phenomenon of 
isomerism. 

648. The Valence of Carbon. — If we consider the six com- 
pounds of carbon, CH 4 , C 2 H 2 , C 2 H 4 , C 2 H 6 , C 3 H 6 , and C 3 H 8 (see 
Table IV, 63), we shall have much difficulty in deciding upon 
the valence of carbon. In CH 4 , assuming the valence of hydro- 
gen to be one, carbon evidently has a valence of four, but this 
value does not seem to harmonize with the formulae of the other 
substances. Consistent conclusions are to be reached only by 
considering the question of the graphic formulae of the substances 
(323). If we start with the plausible assumption that carbon 
has a valence of four, while hydrogen and chlorine each have 
unit valence, we may write, as the graphic formulae of methane 
and the substances formed from it by the action of chlorine, the 
following : 

H H H H CI 

I I I I I 

H— C— H H— C— CI H— C— CI CI— C— CI CI— C— CI 

I I I I 

H H CI CI CI 

Since we write for water the graphic formula 

H-O-H, (323) 



406 Introduction to General Chemistry 

thereby assuming the valence of oxygen to be two, we may 
write for methyl alcohol, CH 4 (645), the graphic formula 

H 

H— C— O— H 

I 
H 

It seems probable that these graphic formulae show the actual 

relations of the atoms to one another in the molecule. If so, 

they show the structure of the molecule, and in consequence 

they may be called structural formulae. As stated earlier (323), 

the lines joining the symbols in such formulae are called bonds. 

Only one possible arrangement of two atoms of carbon and 

six atoms of hydrogen (C 2 H 6 ) satisfies the condition that the 

valence of carbon is four and that of hydrogen one, namely the 

following : 

H H 

I I 
H— C— C— H 

I I 
H H 

This is therefore the accepted formula of ethane (643). 

649. The Structural Formulae of Ethyl Alcohol and Methyl 
Ether. — Ethyl alcohol and methyl ether are both represented 
by the simple formula C 2 H 6 0. There are two structural 
possibilities: 

H H H H 

I I I I 

H— C— C— 0— H and H— C— O— C— H 

I I I I 

H H H H 

Since methyl alcohol is 

H 

I 
H— C— O— H 

I 
H 

and since methyl and ethyl alcohol show similar chemical 
behavior, we may assume provisionally that the first of the 



Carbon and Carbon Compounds 407 

preceding formulae is that of ethyl alcohol and the second that 
of methyl ether. 

The formation of ethyl ether from ethyl alcohol could then 
be considered to take place by the removal of a molecule of 
water from each pair of alcohol molecules, thus : . 

H H H H 

I I 1 : i I I 

H— C— C— 0-4-H+H— O-hC— C— H 

I I ! ! I I 

H H H H 

which leads to the following as the formula for ethyl ether, 

C 4 H I0 O: 

H H H H 

I I I I 

H— C— C— O— C— C— H 

I I I I 

H H H H 

This conclusion is supported by the fact that this formula is 
entirely analogous to that assigned to methyl ether. In general, 
the structure of the molecules of any substance is discovered 
by a study of its reactions and mode of synthesis. 

650. The Cause of Isomerism. — Isomerism finds a simple 
explanation in the assumption that the atoms composing mole- 
cules of isomers are arranged in two (or more) ways, each in 
harmony with simple valence laws. Hundreds of cases of 
isomerism are all adequately explained in this fashion. 

651. Discussion. — We have now seen how a number of the 
simpler organic compounds can be obtained from natural sources. 
We have also learned a little regarding the reactions of organic 
substances. 

The disclosure of the isomerism of ethyl alcohol and methyl 
ether demanded an explanation which was found in the assump- 
tion that the atoms of a molecule are joined to one another in 
definite fashion and always in accord with simple fixed rules of 
valence. 

The importance of structural formulae to the chemist can 
well be illustrated with the examples of methyl ether and ethyl 



408 Introduction to General Chemistry 

alcohol. The OH group attached to a carbon atom which carries 
no other oxygen atom is called an alcohol group. When an 
organic chemist sees this group in a structure he knows that the 
substance must have the properties of- an alcohol modified to 
some degree, according to what other groups are near it in the 
molecule in question. An oxygen atom joined on either hand 
to a carbon atom which carry no other oxygen atoms is the 
group characteristic of the oxygen ethers. The latter are very 
different substances from alcohols and have their own character- 
istics. Thus at a glance the trained chemist can read the char- 
acteristics of a substance by observing the arrangements of the 
atoms in the structure formula of the molecule and recognizing 
groups which have pronounced characteristics. 

Most of the substances encountered in preceding chapters 
belong to one or another of five principal classes : elements, oxides, 
acids, bases, or salts. Among organic compounds we find a large 
number of new and important classes. The paraffine hydro- 
carbons, C M H 2M+2 , starting with methane form the simplest class 
of organic substances. The halogen derivatives like methyl 
chloride, CH 3 C1, and chloroform, CHC1 3 , form another class. 
The alcohols and the ethers are also important classes. The 
balance of this chapter will be devoted to the brief description 
of several other classes of carbon compounds. It is necessary 
that the student should have at least a slight acquaintance 
with the more important classes of organic substances if he 
wishes to get an insight into the nature of foodstuffs, as well as 
of those organic substances which play so important a role in 
our modern daily life. Among the latter are found medicinals, 
dyes, explosives, photographic developers, perfumes, poisons, etc. 

It is not to be expected that the student who reads, however 
carefully, the balance of this chapter will get very definite ideas 
of the methods of preparation and properties of the various 
classes of substances there described. Familiarity with organic 
substances can be gained only by prolonged and detailed study 
in both text and laboratory. Our object in presenting the topics 
about to be discussed will be largely attained if we impress upon 
the student that carbon forms an enormous variety of com- 



Carbon and Carbon Compounds 409 

pounds, that these compounds are of known molecular structure, 
that they fall into definite classes according to the presence in 
them of certain active groups of atoms, and that the members 
of each class show similar behavior, namely, that of the active 
groups, the characteristics of each of which may be modified 
somewhat according to the nature of other groups present in 
the same molecule. 

In what follows, bonds between characteristic groups will 
frequently be represented by dots or entirely omitted. Thus 
CH 3 - OH will be written CH 3 • OH, or for still greater simplicity 
CH3OH. 

652. Aldehydes. — The mild oxidation of ethyl alcohol gives 
acetaldehyde. The reaction is best carried out with a mixture 
of a dichromate and sulfuric acid. Omitting details, the equa- 
tion is 

CH 3 • CH 2 • OH+0->CH 3 • CO • H+H 2 0. 

Acetaldehyde is a very volatile liquid. 

The oxidation of methyl alcohol, CH 3 OH, gives the closely 
related substance formaldehyde, H • CO • H. This is a gas of 
pungent odor which comes on the market as a 40 per cent solu- 
tion in water called formalin. It is extensively used as a 
germicide and antiseptic and also in the manufacture of other 
important organic substances. In general, the oxidation, of any 
alcohol of the formula R • CH 2 • OH, where R is H or a hydro- 
carbon radical, as for example C 2 H S or C 3 H 7 , gives an aldehyde, 
R • CO • H. The aldehydes are good reducing agents because 
they are readily oxidized to the corresponding acids. For 
example, acetaldehyde may be oxidized by suitable agents to 
acetic acid: 

CH 3 • CO • H+0->CH 3 • CO • OH. 

653. Acetic Acid, HC 2 H 3 2 . — As stated earlier (157), acetic 
acid is the most important component of vinegar, of which it 
forms about 4 per cent. Cider vinegar is made from cider, the 
juice of apples. The sweetness of fresh cider is due to glucose. 
Upon standing several days the glucose gradually ferments, 
owing to the growth of yeast, the germs of which are always 



410 Introduction to General Chemistry 

present in the dust of the air, the glucose changing to alcohol and 
C0 2 . The fermented product is popularly known as hard cider. 
If hard cider is allowed to stand in an open or loosely stoppered 
vessel it changes in a few weeks into vinegar. This change is 
the result of the oxidation of alcohol to acetic acid : 

C 2 H 6 0+0 2 ^>HC 2 H 3 2 +H 2 0. 

The oxidation requires the assistance of a catalytic agent pro- 
duced by the so-called vinegar plant, or mother of vinegar 
(micoderma aceti). "White distilled vinegar" is made from 
dilute alcohol produced from corn, substantially in the manner 
already described (640) . The dilute alcohol is allowed to trickle 
slowly through large casks filled with beechwood shavings, 
coated with the slimy mother of vinegar, while oxygen is 
furnished by a countercurrent of air that enters near the bot- 
tom and passes out at the top of the cask. This so-called 
"quick process" produces finished vinegar in eight to ten days. 

Pure acetic acid is a colorless liquid having a sharp, char- 
acteristic odor. When free from water it solidifies at 17 to 
glassy crystals called glacial acetic acid. We have already 
studied the reactions of solutions of acetic acid and its salts 
(157,424,456). 

654. The Graphic Formula of Acetic Acid. — Expressed graphi- 
cally, we have as the equation for the oxidation of alcohol, 

H H H 

I I ' I 

H— C— C— O— H+0 2 ^H— C— C— O— H+H 2 

I I I II 

H H HO 

The group -C-OH, or briefly -CO • OH, is the carboxyl 

II 
O 

radical. One of the oxygen atoms of acetic acid is said to be 
attached to one of the carbon atoms by a double bond (324). 
Of the four H atoms of acetic acid, one occupies a unique posi- 
tion in that it is attached to an atom of O, while the others are 
attached to one of the carbon atoms. Since only one of the H 



Carbon and Carbon Compounds 411 

atoms of the molecule is ionizable we may safely conclude that 
it is the one attached to oxygen. 

655. The Fatty Acids. — There is a series of acids closely 
related to acetic acid, all having a carboxyl radical attached 
to a hydrocarbon radical. The general formula of such acids 
is R • CO • OH, where R stands for hydrogen or any hydro- 
carbon radical such as C 2 H S , C 3 H 7 , C 4 H 9 , etc. The names and 
formulae of a few of the more important are given in Table XXX. 

TABLE XXX 

Formic acid H • CO • OH. 

Acetic acid CH 3 • CO • OH 

Propionic acid C 2 H S • CO • OH 

Butyric acid C 3 H 7 • CO • OH 

Palmitic acid C IS H 3I • CO • OH 

Stearic acid C I7 H 3S • CO • OH 

These acids are known as fatty acids, because some of them are 
obtained from fats (680). 

656. Ketones. — Calcium acetate, Ca(C 2 H 3 2 ) 2 , is an impor- 
tant article of commerce made by the action of acetic acid on 
limestone (CaC0 3 ). When calcium acetate is strongly heated 
it decomposes into CaC0 3 and acetone, C 3 H60: 

Ca(C 2 H 3 2 ) 2 -> CaC0 3 + C 3 H 6 0. 

The acetone distils off and is condensed to a liquid. Purified 
acetone is a colorless liquid of mild but peculiar odor. It boils 
at 56 . It mixes with water or alcohol in all proportions. It is 
an excellent solvent for many organic substances, and it is also 
used in the preparation of several important organic compounds, 
of which chloroform is one. The structural formula of acetone 
is indicated by its formation from calcium acetate: 



H H H H 

I ? 

H— C— C^O— Ca— O— C 

I II ! II 

HO! O 



I I I AK 

C— H->H— C— C— C— H + Ca< >C = O 

I I II I xK 

H H O H 



412 Introduction to General Chemistry 

Acetone is the simplest member of a class of substances called 
ketones, The general formula for a ketone is R x • CO • R 2 , 
where R x and R 2 represent the formulae of hydrocarbon radicals. 
In acetone R z and R 2 are both methyl, CH 3 ; but if, say, R x is 
ethyl, C 2 H S , and R 2 is propyl, C 3 H 7 , the formula of the ketone 
would be C 2 H 5 • CO • C 3 H 7 . 

Ketones are closely related to aldehydes, since if R 2 is H we 
have R r • CO • H, the formula of an aldehyde. They are 
reducing agents but are not as active as aldehydes. 

657. Esters. — The esters are an important class of com- 
pounds, inasmuch as all animal and vegetable fats and oils are 
included therein. One of the simplest esters, ethyl acetate, 
CH 3 COOC 2 H 5 , is obtained by the action of ethyl alcohol on 
acetic acid: 

CH 3 • CO • OH+C 2 H 5 • OH->CH 3 ■ CO • OC 2 H 5 +H 2 0. 

Ethyl acetate is a colorless liquid, boiling at 75 . It has a 
rather pleasant odor and is somewhat soluble in water. 

Just as most acids can form salts with most bases, so most 
acids can form esters with most alcohols. However, alcohols 
are not bases, since they do not yield OH - ions, and esters are 
not salts. Their water solutions are not ionized, and they do 
not give the ionic reactions shown by solutions of the acids from 
which they are derived. The formation of esters requires in 
general the stimulus of H + ions as a catalytic agent. In mak- 
ing ethyl acetate we add to the mixture of acetic acid and alcohol 
some HC1 or H 2 S0 4 . The union of acetic acid and alcohol does 
not take place completely but reaches equilibrium when about 
two-thirds of the possible amount of ester has been formed. 
This is because the reaction is reversible. 

The change of ester and water into acid and alcohol is pro- 
moted by the presence of much water and also by the catalytic 
influence of acids. The speed of ester formation and also the 
speed of reaction of ester and water are greatly increased with 
increase of temperature. 

Methyl acetate, CH 3 CO • OCH 3 , closely resembles ethyl 
acetate. It is miscible with water in all proportions. Other 



Carbon and Carbon Compounds 413 

esters, including fats and oils, will be considered in the next 
chapter. 

658. Amines. — The amines are derivatives of ammonia and, 
like the latter, are base-forming substances capable of yielding 
salts with acids. The simplest member of the class is methyl 
amine, CH 3 NH 2 . If methyl iodide (660) and ammonia are 
mixed they unite to form methyl ammonium iodide: 

CH3I+NH3 -> CH3NH3I. 

The product is a soluble salt of the base CH 3 NH 3 OH. This 

unstable base, which is set free by the action of sodium hydroxide 

on the salt, easily dissociates into methyl amine and water. 

Methyl amine is a colorless gas with an odor resembling ammonia. 

It is abundantly soluble in water, with which it partially unites, 

thus: 

CH 3 - NH 2 +H 2 O^CH 3 • NH 3 • OH. 

Ethyl iodide, C 2 H 5 I, and ammonia give ethyl ammonium 
iodide, C 2 H S • NH 3 • I, from which we readily obtain ethyl 
amine, C 2 H 5 • NH 2 , a substance closely resembling methyl 
amine. 

Methyl amine acts on methyl iodide as follows: 

CH 3 • NH 2 +CH 3 I-»(CH 3 ) 2 NH 2 I, 

from which dimethyl amine, (CH 3 ) 2 NH, is obtained by the 
action of alkali. Dimethyl amine, by further action of methyl 
iodide, yields (CH 3 ) 3 NHI, from which by the action of alkali 
we get trimethyl amine, (CH 3 ) 3 N (59). Amines of various kinds 
are usually found among the products of decomposition of 
proteins (685). Trimethyl amine, for example, is contained in 
herring brine. 

659. Amides. — The interaction of ethyl acetate and ammonia 
gives acetamide and alcohol: 

CH 3 COOC 2 H s +NH3->CH 3 CONH 2 +C 2 HsOH. 

Acetamide is also easily made by distilling ammonium acetate: 
CH 3 COONH 4 -> CH 3 CONH 2 +H 2 0. 



414 Introduction to General Chemistry 

The substance is a white crystalline solid easily soluble in water. 
It unites with hydrochloric acid to form a saltlike compound, 
CH3CONH3CL This fact shows that the NH 2 radical in an 
amide has still some basic properties. 

Most organic acids and some mineral acids are able to form 
amides. Carbonic acid, for example, forms the amide C0(NH 2 ) 2 . 
This important substance is commonly known as urea. By far 
the larger part of the nitrogen content of the food of all animals 
is excreted as urea. Urea reacts slowly with water to form 
ammonia and carbon dioxide: 

CO(NH 2 ) 2 +H 2 ^ 2NH3+ C0 2 . 

660. Ethylene, C 2 H 4 . — A modification of the process of mak- 
ing ether (642) yields ethylene. To make ethylene, a mixture 
of alcohol with six parts by weight of concentrated sulfuric acid 
is heated to 165 , and a mixture of one part of alcohol to two 
parts of sulfuric acid is dropped in slowly. The gas C 2 H 4 is 
given off. The reactions in this case are probably also complex, 
but the net result is the decomposition of alcohol into water and , 
ethylene : 

C 2 H 6 0->H 2 0+C 2 H 4 . 

Ethylene is a colorless gas nearly insoluble in water. It burns 
with a luminous flame. It gives several interesting and impor- 
tant reactions. With Cl 2 it unites to form ethylene chloride; 

C 2 H 4 -f- Cl 2 -> C 2 H 4 C1 2 , 

a colorless, oil-like liquid, boiling at 84 . Ethylene chloride is 
insoluble in water. It has none of the properties of a salt. 
Ethylene also unites with bromine, thus: 

C 2 H 4 +Br 2 ->C 2 H 4 Br 2 . 

The product, ethylene bromide, resembles the chloride. Eth- 
ylene and HBr unite readily to form C 2 H 5 Br, ethyl bromide, a 
liquid boiling at 3 8°: 

C 2 H 4 +HBr->C 2 H 5 Br. 



Carbon and Carbon Compounds 415 

Ethylene and HI give ethyl iodide, boiling-point 72°: 
C 2 H 4 -f-HI->C 2 H 5 I. 

Neither ethyl bromide or ethyl iodide has any of the properties 
of a salt. Methyl iodide, CH 3 I, closely resembles ethyl iodide. 
It is used in the preparation of other organic compounds (658). 

661. The Structural Formula of Ethylene. — The fact that 
ethylene unites with chlorine, hydrobromic acid, etc., leads to 
the conclusion that two of the valence bonds of the carbon 
atoms of the C 2 H 4 molecule are either free, 

H H 

I I 
H— C— C— H, 

I I 

or, more probably, are attached to, or satisfied by, one another 
so as to form a double bond (324) between the two carbon atoms: 

H H 

I I 
H— C=C— H. 

In the reaction with chlorine, for example, the extra bonds unite 

with chlorine to give 

H H 

I I 
H— C— C— H. 

I I 
CI CI 

662. The Ethylene Series, C W H 2M . — A long series of hydro- 
carbons which have twice as many H as C atoms per molecule 
and of which the first member is C 2 H 4 is known. The second 
member is propylene, CH 3 — CH= CH 2 . There are theoretically 
just three butylenes, C 4 H 8 : 

(1) CH 3 • CH 2 • CH = CH 2 

(2) CH 3 . CH = CH- CH 3 

CH 3 v 

(3) >C = CH.CH 3 
CH/ 



4i 6 Introduction to General Chemistry 

These are all known. All of the members of this series unite 
with Cl 2 , Br 2 , HI, etc. For example, 

CH 3 - CH 2 - CH = CH 2 + Cl 2 -> CH 3 - CH 2 - CHC1 - CH 2 C1. 

663. The Acetylene Series, C„H2k- 2 - — The acetylene series 
of hydrocarbons has the general formula C M H 2W _ 2 . It is headed 
by acetylene, C 2 H 2 (49, 83). It is probable that the four extra 
valence bonds of acetylene, 

H— C— C— H, 

I I 

are united in pairs, so that the carbon atoms are joined by a 
triple bond, 

H— CeeeC— H. 

In accord with this view we should expect acetylene to unite 

with chlorine, thus: 

CI CI 

I I 
H— C=C— H+ 2 Cl 2 - H— C— C— H 

I I 
CI CI 

and, in fact, it actually behaves in this way. Because of the 
fact that C 2 H 4 and C 2 H 2 have unsatisfied valences, as shown 
by their union with Cl 2 , HBr, etc., these hydrocarbons are said 
to be unsaturated. 

664. Isomerism of Hydrocarbons. — All of the compounds 
C„H 2M+2 , where n is four or more, can exist in two or more 
isomeric forms. The simplest case is that of the butanes, 
C 4 H I0 , of which two are known: normal butane, 

CH 3 — CH 2 — CH 2 — CH 3 , 

and isobutane, 

CH 3V 

>CH-CH 3 . 
CH/ 



Carbon and Carbon Compounds 417 

There are three isomeric pentanes : 

. (1) CH 3 • CH 2 - CH 2 - CH 2 • CH 3 

CH 3 v 

(2) >CH-CH 2 -CH 3 
CH/ 

CH 3 

I 

(3) CH-C-CH 3 

I 
CH 3 

The first of these is said to have a straight carbon chain; the 
second and third have branched chains. The B.P. and M.P. 
columns of Table XXVIII (643) refer to the normal hydro- 
carbons, with straight chains. 

665. Some Common Organic Acids. — We shall now briefly 
describe several of the commoner organic acids. The simplest 
fatty acid is formic acid, HCOOH. It is present in the bodies 
of ants and constitutes the poison of bees' stings. It can be 
made artificially in several ways. It is a colorless liquid dis- 
solving .easily in water to form a moderately well-ionized acid 
solution. 

Oxalic acid, H 2 C 2 4 , is a white crystalline substance. Its 

structural formula is 

COOH 

I 
COOH. 

It is a dibasic acid (102), forming both acid and neutral salts, 
as for example NaHC 2 4 and Na 2 C 2 4 . The free acid, which is 
easily soluble in water and rather highly ionized, crystallizes 
from water as a hydrate of the formula H 2 C 2 4 • 2H 2 0. The 
acid is decidedly poisonous, probably because of the ease with 
which it decomposes into the powerful poison carbon monoxide 
(632) and water: 

H 2 C 2 4 ->H 2 0+2CO. 

This decomposition takes place rapidly when the acid is heated 
with concentrated sulfuric acid, and so affords a good way of 
making carbon monoxide. Oxalic acid occurs in some plants. 



4i 8 Introduction to General Chemistry 

The chief acid present in sour milk is lactic acid, the graphic 
formula of which is CH 3 • CHOH • COOH. It gives a colorless 
water solution of a pleasant sour taste. Tartaric acid, a dibasic 
acid having the formula H 2 C 4 H 4 6 , is abundant in grapes in the 
form of its acid potassium salt, KHC 4 H 4 6 . This salt is obtained 
as argol in large amount in the manufacture of wine. The 
refined salt is known as cream of tartar. High-grade baking 
powder is a mixture of cream of tartar and sodium bicarbonate. 
The dry mixture is fairly stable. In the presence of water the 
substances react thus: 

KHC 4 H 4 6 +NaHC0 3 ->KNaC 4 H 4 6 +C0 2 +H 2 0. 

The leavening power of baking powder is due to the C0 2 given 

off (cf. 593). Tartaric acid forms colorless crystals easily soluble 

in water. Its graphic formula is 

O 

COOH C • OH 

I I 

CHOH H • C • OH 

I or I 

CHOH H - C • OH 

I I 

COOH C • OH 

II 

o 

Citric acid, H 3 • C 6 H s 7 , is the acid of lemon and other 
citrus fruits. It is a tribasic acid (159) rather closely resembling 
tartaric acid but having a more complex formula. Prussic or 
hydrocyanic acid, HNC, is an extremely poisonous and very 
weak acid. Its salts, sodium cyanide, NaNC, and potassium 
cyanide, KNC, are made from calcium cyanamid (526). Both 
salts react readily with either oxygen or sulfur, forming the 
corresponding cyanate or sulfocyanate respectively: 

2 KNC+0 2 ->2KNCO, 
KNC+S^KNCS. 

The cyanide ion unites with many metal ions to form com- 
plex cyanide ions. Salts containing these ions, such as potas- 



Carbon and Carbon Compounds 419 

sium ferrocyanide, K 4 Fe(CN) 6 , and potassium ferricyanide, 
K 3 Fe(CN)6, are important analytical reagents. 

666. The Aliphatic and Aromatic Series of Organic Com- 
pounds. — Broadly speaking, organic compounds constitute two 
great series, the aliphatic and the aromatic. The substances 
thus far mentioned are all members of the aliphatic series. We 
may consider that the paraffin hydrocarbons, C w H 2w+2 , are 
the fundamental substances from which all aliphatic compounds 
are derived. Thus paraffin hydrocarbons by loss of two 
hydrogen atoms per molecule give members of the ethylene 
series, C M H 2W , or by loss of four hydrogen atoms per molecule 

TABLE XXXI 

Formula of Typical 
Class of Compound Compound 



Hydrocarbon CH 3 

Halide CH 3 

Alcohol CH, 

Aldehyde CH 3 

Acid CH 3 

Ester CH 3 

Amide CH 3 

Ether CH 3 

Ketone CH 3 

Amine CH 3 



H 

CI 

OH 

CO • H 

CO • OH 

CO • OCH 

CO • NH 2 

0-CH 3 

CO • CH 3 

NH 2 



give members of the acetylene series. Substitution of hydrogen 
by halogens or by various radicals such as methyl, CH 3 ; 
hydroxyl, OH; carboxyl, COOH; amid, NH 2 ; etc., gives rise 
to the various classes of compounds already briefly studied. 

Table XXXI shows the formulae of typical members of the 
most important classes of aliphatic compounds. 

Instead of the methyl radical, CH 3 (Table XXXI), we may 
have ethyl, C 2 H S , propyl, C 3 H 7 , or any radical derived from any 
hydrocarbon in each case and thus obtain a very great variety 
of substances. All of these substances belong to the aliphatic 
series. 

The aromatic compounds differ from the aliphatic in that 
they are derived from hydrocarbons entirely different from the 
paraffins. The most fundamental and characteristic aromatic 
hydrocarbon is benzene. 



420 Introduction to General Chemistry 

667. Benzene, C 6 H 6 . — Benzene is a colorless liquid, boiling 

at 79 and practically insoluble in water. It is a by-product of 

the manufacture of coke and coal gas (634). When coal is 

heated in the absence of air it yields, in addition to coke and 

gas, a large amount of black liquid tar. By the distillation of 

tar a number of very important aromatic hydrocarbons are 

obtained. One of the most useful of these is benzene. We 

shall give the structural formula of benzene without attempting 

to justify our reasons therefor, since that story (although a 

most interesting one) is entirely too long for this text. This 

formula is 

H 

/ C \ 
HC CH 

HC CH 

H 

It represents a ring of six carbon atoms, with three single and 
three double bonds, and six hydrogen atoms, one attached to 
each atom of carbon. 

Benzene differs from the paramne hydrocarbons (which are 
very inactive) in being remarkably active chemically. This 
activity is shown in two ways: (1) by substitution of various 
radicals for one or more hydrogen atoms of each molecule, and 
(2) by addition after the manner of unsaturated compounds 
(660) but much less readily. 

668. Other Aromatic Hydrocarbons. — Next to benzene the 
simplest aromatic hydrocarbon is toluene, or methyl benzene, 
C 6 H S CH 3 , the structural formula of which is 

H 

HC CCH 3 

II I 

HC CH 

H 



Carbon and Carbon Compounds 421 

In order to save labor the C 6 H 5 radical is frequently represented 
graphically by a hexagon called the benzene ring, so that the 
formula of toluene is written thus: 



n 



CH 3 

Toluene is a liquid, boiling at no° and closely resembling 
benzene. It is obtained from coal tar. Xylene, also obtained 
from the same source, is dimethyl benzene, C 6 H 4 (CH 3 ) 2 . In 
writing the structural formula of xylene we note that there are 
three possible arrangements: 

CH3 CH 3 

Am. /\ /\ 



f 



\/ 



CH, 



\s 



CH 3 

CH 3 



J 



As a matter of fact three different xylenes are known. These 
are called ortho, meta, and para xylene, respectively written 
o-xylene, m-xylene, and p-xylene. 

669. Naphthalene, C I0 H 8 . — Naphthalene is a white crystalline 
substance obtained from the high-boiling portion of coal tar. 
It is extensively used in the household under the name of 
moth balls. Its structure is represented thus: 

H H 

/ C \/ C \ 
HC C CH 

HC C CH 

H H 

As great a variety of organic substances are derived from 
naphthalene as from benzene. For example, we have two 
methyl naphthalenes, C I0 H 7 • CH 3 : 

CH 3 

ill II I™ 3 



422 Introduction to General Chemistry 

The first is called alpha and the second beta methyl naphthalene. 
Naphthalene and its derivatives are important starting materials 
for the manufacture of dyestuffs. 

670. Aromatic Alcohols and Aldehydes. — Toluene, C 6 H S • 
CH 3 , can be considered as derived from methane, CH 4 , by 
the substitution of the phenyl radical, C 6 H 5 , for one hydro- 
gen atom of CH 4 . The substitution of C6H 5 for H in methyl 
alcohol, CH 3 • OH, would give C6H S • CH 2 • OH. The sub- 
stance actually exists and is known as benzyl alcohol. Like 
other alcohols, it can be oxidized to an aldehyde called benzal- 

dehyde: 

C 6 H S • CH 2 • OH+0 2 ->C 6 H 5 • CO - H+H 2 0. 

This aldehyde is identical with the principal constituent of 
the oil of bitter almonds. Benzaldehyde can be made from 
toluene by converting the latter into a chlorine compound and 
then treating this product with water: 

C 6 H 5 • CH 3 +2C1 2 ->C 6 H S - CHC1 2 +2HC1, 
C 6 H S • CHC1 2 +H 2 0->C 6 H S . CO • H+2HCI. 

Benzaldehyde is slowly oxidized by contact with air to form 
benzoic acid, C6H 5 • CO • OH: 

2 C 6 H S • CO • H+0 2 -> 2 C 6 H 5 • CO • OH. 

671. Benzoic Acid, C 6 H 5 • CO • OH. — This important acid 
is a white crystalline solid which can be made in several ways in 
addition to the one just mentioned: It is used extensively in 
the form of its sodium salt, sodium benzoate, C 6 H 5 CO • ONa, 
as a preservative for catsup and other articles of food. It can 
be used legally in the United States as- a food preservative if its 
presence is indicated on the label. It occurs naturally as a 
constituent of cranberries. 

Benzoic acid forms salts with bases of all kinds. With 
alcohols it forms esters. The latter are fragrant liquids. Ethyl 
benzoate, C6H S • CO • OC 2 H 5 , is a colorless liquid, boiling at 211 . 

672. Phenol, C 6 H 5 • OH. — Phenol (popularly known as 
carbolic acid) is contained in coal tar, from which it is separated 
in crude form by distillation. Phenol is also made on a large 



Carbon and Carbon Compounds 423 

scale from benzene. The latter substance reacts slowly with 
sulfuric acid, forming benzene sulfonic acid and water, thus: 

C 6 H 6 +H 2 S0 4 ->C 6 H 5 S0 2 OH+H 2 0. 

The product is an acid the sodium salt of which when fused with 
sodium hydroxide gives phenol: 

C 6 H s S0 2 ONa+NaOH .-> C 6 H 5 OH+Na 2 S0 3 . 

Phenol is a white crystalline substance having a peculiar, char- 
acteristic odor. It is moderately soluble in water. It is a 
violent poison and is extensively used as a germicide. Phenol 
is a very weak acid and forms salts with strong bases : 

C 6 H 5 OH+NaOH -> C 6 H 5 ONa+H 2 0. 

It is interesting to contrast phenol and ethyl alcohol. The 

formulae of the two substances 

H 



H- 



show that both contain the hydroxyl radical united to a hydro- 
carbon radical. In consequence we might expect similar 
properties, but we find quite the contrary. Phenol shows but 
few of the characteristic chemical properties of an alcohol. 

673. Aromatic Nitro Compounds. — Aromatic hydrocarbons, 
like benzene, react very readily with concentrated nitric acid 
in a peculiar way, as illustrated by the following equation: 

C 6 H 6 +HN0 3 -> C 6 H 5 N0 2 + H 2 0. 

The new product is nitro benzene, a light-yellow colored liquid 
of aromatic odor. It is not soluble in water and is not a salt. 
The structural formula is 

0« 



H H 


/ C \ 


1 | 


HC COH 


C— C— OH 


II 1 


i 1 


HC CH 


H H 


\S 




H 



424 Introduction to General Chemistry 

Toluene and nitric acid give two isomeric nitro compounds, 
CH 3 ■ C 6 H 4 • N0 2 . These have the following formulae: 

CH 3 CH 3 

M N ° 2 M 

\/ \/ 

N0 2 

The left-hand formula is that of ortho nitro toluene ; the right- 
hand one that of para nitro toluene. A third nitro toluene can 
be made by indirect methods. This substance, called meta 
nitro toluene, has the formula 



r 



CH 3 
\ 



\/ 



N0 2 



The further action of nitric acid on either ortho or para nitro 
toluene gives finally tri nitro. toluene, CH 3 • C6H 2 • (N0 2 ) 3 . 
This substance is the violent explosive so extensively used in 
the war and known popularly as T.N.T. Its formula is 

CH 3 
2 N /\ N0 2 



N0 2 

The action of nitric acid on phenol gives first a mixture of 
ortho and para nitro phenol, 

OH 0H 

|^1 NO, 

x ' NO 



Further action of nitric acid finally yields tri nitro phenol, or 
picric acid, C 6 H 2 (N0 3 ) 3 OH. The nitro phenols are much 
stronger acids than phenol itself. In fact, picric acid is nearly 
as strong (highly ionized) an acid as hydrochloric. 

Ammonium picrate, C 6 H 2 (N0 3 ) 3 ONH 4 , is a powerful explo- 
sive. It has been extensivelv used in the war. 



Carbon and Carbon Compounds 425 

674. Aromatic Amines. — Nitro compounds are easily acted 
on by reducing agents, as illustrated by the case of nitro benzene : 

C 6 H 5 • N0 2 + 3 H 2 ->C 6 H 5 • NH 2 + 2 H 2 0. 

The new product is called aniline. It is the simplest aromatic 
amine. The aromatic amines resemble the aliphatic amines 
(658). They are base-forming substances and are to be con- 
sidered as ammonia in which hydrogen has been replaced by an 
aromatic radical. Aniline unites with hydrochloric acid to 
form a true salt, a chloride: 

C 6 H 5 NH 2 +HC1 -> C 6 H 5 NH 3 C1. 

This salt corresponds to NH 4 C1. A great variety of aromatic 
amines can be made by the reduction of nitro compounds. 
Usually, the reduction is carried out by mixing the nitro com- 
pound with zinc or iron and hydrochloric acid. The hydrogen 
liberated by the action of the metal and acid then reacts with 
the nitro compound in the way above indicated. Aniline and 
other aromatic amines are made in immense quantities to be 
used as intermediates in the manufacture of so-called aniline 
or coal-tar dyes. Further reference to this subject will be found 
in the next chapter. 



CHAPTER XXVI 
ORGANIC COMPOUNDS. II 

675. Introduction. — We have promised to show the reader 
some of the successes achieved in organic chemistry as a result 
of the systematic study of the science. First we shall take up 
the chemistry of foods and next the chemistry of explosives and 
the related substances, such as celluloid, artificial silk, etc. 
After this we shall treat briefly the poison gases used in war- 
fare, then the synthesis of essential oils, perfumes, spices, dyes, 
medicinals, and rubber. This may seem a bewildering list, but 
organic chemists can say without fear of contradiction that 
they have accomplished a great deal more than is even casually 
mentioned in this chapter. Finally we shall discuss briefly the 
chief sources of materials for the manufacture of organic chem- 
icals. 

676. The Three Classes of Foods. — All foods fall into three 
great classes: the fats, including also edible oils; the carbo- 
hydrates, comprising starches and sugars; and the proteins, in 
which class are included eggs, lean meat, and certain nitrogenous 
vegetable products. A well-balanced diet for man should be 
made up of foods of all three of these classes. Fats and carbo- 
hydrates are compounds of carbon, hydrogen, and oxygen only; 
while all proteins, in addition to these three elements, contain 
also nitrogen. Of the three classes the fats are from the chemical 
point of view the simplest, and their chemistry was worked out 
long before that of the other two classes. The chemistry of the 
carbohydrates was well cleared up during the last two decades 
of the nineteenth century. The chemistry of the proteins is 
far more complex and is even today far from completely solved. 

677. Fats. — The term fats includes liquid as well as solid 
animal and vegetable products. Liquid fats like olive oil, cotton- 
seed oil, peanut oil, castor oil, and linseed oil are chemically 
very different from the paraffin or mineral oils described in the 

426 



Organic Compounds 427 

foregoing chapter. Most natural fats (butter fat for example) 
are mixtures of several chemical compounds all of which belong 
to a single group of organic substances, the esters (657). The 
chemical nature of these esters is most readily shown by the 
conversion of fats into soap. 

678. Soap. — Any fat is changed into soap when it is boiled 
with a solution of sodium hydroxide. If a good grade of white 
soap is dissolved in water and the solution acidified with hydro- 
chloric acid a dense white precipitate forms. The evaporated 
filtrates yield only common salt. The white precipitate is a 
mixture of three or four fatty acids (655). Among the com- 
monest fatty acids obtained in this way are palmitic acid, 
C I5 H 3I • CO • OH, and stearic acid, C I7 H 35 • CO • OH. 

The radicals C I5 H 3I and C I7 H 35 form straight or unbranched 
carbon chains, as illustrated in the following formula for palmitic 

acid: 

HHHHHHHHHHHHHHH 
HC-C-C-C-C-C-C-C-C-C-C-C-C-C-C-CO-OH 
HHHHHHHHHHHHHHH 

Stearic acid contains one more CH 2 group per molecule than 
palmitic acid. 

Soaps are the sodium (or potassium) salts of fatty acids. 
The action of HC1 on sodium palmitate takes place thus : 

C I5 H 3I COONa+HCl->C IS H 3I COOH+NaCl. 

679. Glycerine and Its Esters. — The action of sodium 
hydroxide on a fat always gives in addition to a soap one other 
product, glycerine. Glycerine is a sweet, sirupy, colorless liquid, 
the structural formula of which is 

CH 2 OH 

I 
CHOH 

I 
CH 2 OH 

Glycerine is an alcohol, but it differs from simple (monatomic) 
alcohols like methyl alcohol, CH 3 OH, and ethyl alcohol. 



428 Introduction to General Chemistry 

CH 3 • CH 2 OH, in having three hydroxyl groups in a molecule. 

It is called a triatomic alcohol. It will be recalled (657) that 

alcohols are not bases, in spite of the presence of hydroxyl 

groups. They do not yield OH~ ions. 

Just as acetic acid and ethyl alcohol unite to form ethyl 

acetate (an ester) and water (657), so a fatty acid and glycerine 

can unite to form an ester in which three molecules of the acid 

are combined with one of glycerine. Thus palmitic glycerine 

ester, or palmitin, is 

C IS H 3I COOCH 2 

I 
C I5 H 3I COOCH 

I 
C I5 H 3I COOCH 2 

This substance is one of the principal constituents of beef fat. 
Fats in general are the glycerine esters of various fatty acids. 
Just as ethyl acetate gives, with sodium hydroxide, sodium 
acetate and ethyl alcohol, 

CH 3 • CO • OC 2 H 5 +NaOH -> CH 3 • CO • ONa+C 2 H 5 OH, 

so a fat and sodium hydroxide yield a soap and glycerine. On 
account of the close chemical relation between these two reactions 
the first as well as the second is spoken of as a saponification of 
the ester, although of course sodium acetate is not a soap in the 
ordinary sense of the term. 

680. The Composition of Fats. — We are now in a position to 
understand the cause of the differences between fats from various 
sources. In general, a given sort of fat is a mixture of the 
glycerine esters of several fatty acids. Among such, in addition 
to palmitic and stearic acids, already mentioned, we have oleic 
acid, C I7 H 33 COOH; lauric acid, CnH 2I C00H; caprylic acid, 
C 7 H I5 C00H; caproic aci^CsHxxCOOH; valeric acid,C 4 H 9 C00H; 
and butyric acid, C 3 H 7 COOH. 

Beef fat is composed largely of the esters of palmitic, stearic, 
and oleic acids. These esters are known respectively as pal- 
mitin, stearin, and olein. The first two are solids, while the. 
last is an oil at room temperature. Mutton fat resembles 



Organic Compounds 429 

beef fat in composition but contains a smaller proportion of 
olein, while hog fat in the form of lard contains appreciably 
more olein than beef fat. Butter fat contains, in addition to 
palmitin, stearin, and olein, a considerable proportion of butyrin, 
the glycerine ester of butyric acid. The chief constituent of 
olive oil is olein. The same ester, together with others, com- 
prises cottonseed oil, an edible oil of enormous economic 
importance. 

681. The Hardening of Oils.— Oleic acid, C I7 H 33 COOH, 
differs from stearic acid, Ci 7 H 3S COOH, by two atoms of hydrogen 
per molecule. This difference is the result of one double bond 
(661, 662) between two of the carbon atoms of the C I7 H 33 
radical, which is therefore an unsaturated (663) compound. 
By the addition of hydrogen to the double bond, oleic acid can 
be converted into stearic acid : 

C I7 H 33 COOH+H 2 -> C I7 H 35 COOH. 

By a similar addition of hydrogen, olein is changed into 
stearin. This change is accomplished by the aid of a catalytic 
agent, finely divided metallic nickel. By means of this process 
of hydrogenation, liquid fats like cottonseed oil are readily 
changed into solid or partially solid fats. During the last twenty 
years this so-called hardening of fats has developed into an 
immense industry. The product made from cottonseed oil has 
about the consistency of lard and finds extensive use as a sub- 
stitute for the latter. 

682. The Carbohydrates. — In the preceding chapter starch 
(636) and glucose, or grape sugar (639), were briefly described. 
These substances belong to an important class of organic com- 
pounds known as carbohydrates. This name was chosen because 
these substances are composed of carbon, together with hydrogen 
and oxygen, the two latter in the proportion corresponding 
to water. Thus glucose is C6H I2 06, which is equivalent to 
C6(H 2 0)6- However, the carbohydrates are not simply carbon 
with water of hydration in the sense that Na 2 S0 4 • ioH,0 is 
the hydrate of Na 2 S0 4 . 



43 o Introduction to General Chemistry 

The simplest formula which would represent the composition 
of starch is C6H I0 5 (636); but it is certain that the molecule 
of starch is much larger than that represented by this formula. 
The formula is more correctly written (C 6 H I0 5 ) w , where n is 
an integer probably as large as 30 or 40. Ordinarily the simpler 
formula is employed. 

The hydration of starch to form glucose (638), 

C 6 H I0 O 5 +H 2 O->C 6 H I2 O6, 

is a very important reaction. It takes place readily in acid 
solution by reason of the catalytic action of H ions. The 
higher the temperature of the solution the more rapid the hydra- 
tion proceeds. 

Glucose has a large number of isomers (650), all having, of 
course, the same formula. These sugars, called hexoses, all 
have properties more or less like those of glucose. Le.vulose, 
or fruit sugar, is one of the commonest of the hexoses. It is the 
sugar most abundant in many fruits. 

Ordinary table sugar, commonly known as cane sugar and 
called by chemists sucrose, has the formula C I2 H 22 On. It is 
made from two principal sources, sugar cane and sugar beets. 
It has the same composition in each case. Numerous other 
plants also produce sucrose. Maple sugar is largely sucrose. 

Milk sugar, or lactose, C I2 H 22 0n, is an isomer of cane sugar. 
It is present in cow's milk to the extent of 4 per cent. It is 
much less sweet than cane sugar. Maltose, C I2 H 22 On, is another 
isomer of cane sugar. It is formed by the hydration of starch 
in the presence of a catalytic agent occurring in germinating 
seeds. The reaction may be written 

2C 6 H I0 O 5 +H 2 O -> C I2 H 22 II5 
or better, 

2 (C 6 H I0 5 )„+ «H 2 -> hCmH^Oh. 

The catalytic agent is called diastase. A similar substance, 
ptyalin, is present in saliva. It promotes the digestion of starch 
by hydrolyzing it to maltose. The catalytic agents diastase 
and ptyalin are classed as enzymes. 



Organic Compounds 431 

683. The Structure of the Sugars. — The structure of the 
simpler sugars, like glucose, was worked out at the end of the 
nineteenth century. The following structure of glucose was 
discovered after long experimentation, which established the 
presence of the groups indicated. 

H 2 COH 

I 
HCOH 

I 
HCOH 

HCOH 

I 
HCOH 

I 
HCO 

The five hydroxyl groups behave like those of an alcohol. In 
this respect glucose is an alcohol somewhat resembling glycerine 
(679). One end carbon atom of the glucose molecule forms an 
aldehyde radical (652). The reactions of glucose are those of 
an alcohol and of an aldehyde. Like all aldehydes glucose is 
a good reducing agent. It reduces an alkaline solution of copper 
to cuprous oxide. This reaction, which serves as the best test 
for glucose, is carried out by warming glucose with Fehling's 
solution. This solution is made by mixing copper sulfate solu- 
tion with a solution of sodium tartrate (665) containing an 
excess of sodium hydroxide. In the presence of glucose the 
deep-blue Fehling's solution gives a red precipitate of cuprous 
oxide, Cu 2 0. 

The structure of levulose is represented thus: 

H 2 COH 

I 
HCOH 

I 
HCOH 

I 
HCOH 

I 
CO 

I 

H 2 COH 



432 Introduction to General Chemistry 

It is an alcohol ketone (656). The behavior of cane sugar with 
dilute acids throws much light on its structure, since in this 
reaction it unites with water and forms equal amounts of glucose 
and levulose : 

C12H22OH+H2O -> C6H I2 06-hC6H I2 06. 

From this it follows that in the molecule of cane sugar a molecule 
of glucose is joined with one of levulose, with the elimination of 
a molecule of water. 

Since maltose gives by hydrolysis two molecules of glucose 
its molecule may be considered to be made up of two glucose 
radicals. 

Cane sugar and maltose do not reduce Fehling's solution. 
If their solutions are first hydrolyzed the resulting solution 
reduces Fehling's solution readily. 

684. Cellulose. — Cellulose, which occurs nearly pure in 
cotton, is an isomer of starch (636). Its simplest formula is 
C 6 H I0 O s , but its true formula should be written (C 6 H I0 5 ) w , 
where m is an integer probably even larger than n in the starch 
formula. Cellulose is far less active chemically than starch and 
is practically indigestible by man. It is possible to hydrolyze 
cellulose to glucose, but the reaction takes place slowly. Wood 
and vegetable fiber in general contain a large proportion of 
cellulose. The latter is classed with starch and sugars as a 
carbohydrate. 

685. The Proteins. — We shall use the term protein to include 
the various complex nitrogenous substances forming the charac- 
teristic constituents of lean meat, white of eggs, etc. Most plant 
seeds also contain more or less proteins. Wheat is particularly 
rich in this respect. 

Proteins are complex substances containing in addition to 
carbon, hydrogen, and oxygen a rather large percentage of 
nitrogen and a much smaller percentage of sulfur; a few of the 
proteins also contain phosphorus. Neither the exact formula 
nor the structure of any typical protein is definitely known. The 
composition of albumin (white of egg) is approximately expressed 
by the formula C 720 H II34 N 2I 8S s O 24 8. Although considerable 



Organic Compounds 433 

progress has been made in recent years toward the elucidation 
of the structure of the proteins, much remains to be done. When 
the proteins are heated with acids or alkalies they are split up 
into simpler substances which are found to be nitrogen deriva- 
tives of fatty acids (655). We have seen that acetic acid, 
CH3COOH, for example, forms an amide, CH 3 CONH 2 , acet- 
amide (659). We have also learned something of the amines, 
of which methyl amine, CH 3 NH 2 , is the simplest representative, 
and it will therefore not be surprising to learn that a substance 
having the formula 

H 2 NCH 2 COOH, 

which we may call amino acetic acid, or glycocoll, exists. This 
acid can form an amide, 

H 2 NCH 2 CONH 2 , 

which can unite with one or more molecules of glycocoll to form 
such products as 

H 2 NCH 2 CONHCH 2 CONH 2 
and 

H 2 NCH 2 CONHCH 2 CONHCH 2 CONH 2 . 

Still more complex substances, amino acids, have been built 
up artificially in the laboratory. The fact that these substances 
are identical with, or related to, the decomposition products of 
the proteins leads us to think that the molecules of the latter 
are made up of amino acid radicals, among which are those of 
fatty acids other than acetic. Proteins from different sources 
differ markedly from one another by reason of the kinds and 
relative amounts of the amino acids which compose them. 

Many proteins are of vegetable origin. The cereals like 
wheat, oats, rye, barley, and corn are comparatively rich in 
these nitrogenous substances (511). Beans, peas, and other 
legumes also contain large percentages of proteins, while vege- 
tables contain but small amounts. 

686. Why the Body Needs Food. — The body needs food for 
growth, repair, and the supply of energy. For the adult only 



434 Introduction to General Chemistry 

the last two are of importance. The average amount of carbon 
dioxide exhaled per day by an adult is about 1,100 grams; but 
the amount varies, greatly increasing with the amount of work 
done. Nitrogen is excreted largely as urea (659) but also in 
smaller amounts in the form of other compounds. The average 
daily loss calculated as nitrogen amounts to about 20 grams for 
an adult. Hydrogen and oxygen are also lost in large quantities, 
principally in the form of water and to a smaller extent in com- 
pounds with carbon, nitrogen, etc. These losses must be com- 
pensated by food and water in order to maintain bodily weight; 
for of course the law of the conservation of matter (21) applies 
rigorously to all bodily processes. 

687. The Law of the Conservation of Energy for Bodily 
Processes. — The body expends energy in two ways: in doing 
work and in giving off heat to the surroundings. The source of 
this energy is found in the chemical changes of the food eaten 
and the oxygen inhaled. The amount of energy produced when 
a known amount of a given foodstuff, together with sufficient 
oxygen, is changed to the same products as those formed in the 
body can be determined by means similar to that described 
earlier (357). The energy per gram of a food may be expressed 
in calories and called its fuel value. 

Until recent years it was a question whether the amount of 
energy supplied by the food eaten was exactly equal to that 
expended in the form of work plus that given off as heat when 
the body neither gained nor lost in weight. This important 
problem was solved by the very elaborate experiments of the 
American chemists Atwater and Benedict. These scientists 
constructed a huge calorimeter (357) in which a man could live 
and perform work for hours at a time. The amount of work 
done and heat given off was accurately measured, and the 
energy or calorific value of all food eaten was determined. 
After several years of the most careful work, in which animals 
as well as men were experimented upon, it was conclusively 
proved that the energy given out by the body is exactly equal 
to that produced when the food and oxygen taken are changed 
to the same forms as those excreted by the body. In other 



Organic Compounds 435 

words, it was found that the law of the conservation of energy 
applies rigidly to all bodily processes. 

688. The Science of Dietetics. — The facts set forth in the 
two preceding sections serve as the basis for a scientific treat- 
ment of the subject of nutrition. This branch of science is 
called dietetics. A satisfactory, well-balanced ration must 
supply sufficient amounts of each class of food to compensate 
for the known or determinable body losses of material and at 
the same time yield sufficient energy to enable the performance 
of the required amount of physical work and also keep up bodily 
temperature, all without the loss of body weight. 

The energy requirement of a man doing moderate work is 
about 3,200 kilogram calories. 1 To get this he must consume 
food more than sufficient to satisfy his requirements for carbon ; 
but this will not necessarily also supply his requirements for 
nitrogen. Therefore a sufficient ration will result from eating 
enough protein to compensate the nitrogen loss, and in addition 
enough fats or carbohydrates, or better both, to bring the fuel 
value up to his requirement. The ratio of carbohydrates to 
fat in the diet is not of fundamental importance; furthermore 
the protein consumption may, for some individuals or even for 
whole races, be safely increased far above the necessary minimum. 
For adult male Americans, professional, business men, and 
students, 100 g. of protein, 125 g. of fat, and 400 g. of carbo- 
hydrates, with a total fuel value of 3,200 kilo calories, consti- 
tutes an average daily ration. 

Table XXXII gives the data for sample meals for one day 
for one man doing moderately heavy work. It shows the 
weight in fractions of a pound and in grams of each article, and 
also its protein content and fuel value. 

689. Vitamines. — It would naturally be inferred from the 
discussion of the preceding sections that in providing for a 
dietary sufficient to maintain health and weight it is necessary 
to consider only the protein content and fuel value of the food 

1 On account of the fact that the calorie (in) is so small a unit of heat, a 
unit one thousand fold larger is in common use. This larger unit is the heat 
required to raise the temperature of one kilogram of water one degree. It may be 
called a kilogram caloric, or kilo calorie. 



436 



Introduction to General Chemistry 



supplied. Yet it has long been known that in earlier times 
sailors who lived for long periods on an abundance of food of 



TABLE XXXII 



Article 


Pound 


Percent- 
Grams AGE OF 
Protein 


Protein 
in Grams 


Fuel Value 
in Kilo 
Calories 




Breakfast 


Bread 

Half an orange 

Sugar 

Two eggs 

Butter 

Cup of coffee and cream . . . 


o. 20 
0.30 
0.05 
O. 20 
0.05 


90 

135 

23 

90 

23 


9.2 
0.6 
0.0 

131 
1 .0 


8.28 
0.81 
0.00 
11.79 
0.23 
0.05 


240 

50 

88 

160 

170 

90 










Totals 








21 . 16 


798 












Luncheon 


Tomato soup 

Lamb chops 

Peas 

Stewed apples 

Sugar 

Milk .- 

Bread 

Butter 


0. 20 
O. 20 
0.12 
0. 12 
0.05 
0.50 
0. 20 
0.05 


90 
90 
56 
56 
23 
227 
90 
23 


1 . 1 

13 -5 
8.0 

°-3 
0.0 

3-3 
9.2 
1 .0 


1 .00 
12. 22 

4-50 
0.17 
0.00 

7-49 
8.28 
0.23 


50 

283 

66 

55 
88 

155 
240 
170 


Totals 








33-89 


1,107 












Dinner 




O. 20 
O.40 
O.50 
O. 12 
O.I2 
O.O3 
O. 20 
O.O5 
O. 20 
O. IO 


90 

181 
227 
56 
56 
16 
90 

23 
90 

45 








Roast beef 


150 
1.8 

1 .0 


27.00 
4.09 
0.78 
0.60 


400' 


Baked potato 

Cauliflower \ 

Lettuce 


197 

15 

10 

100 


Bread 

Butter 

Ice-cream 


9.2 
1 .0 
2.6 


8.28 
0.23 
2.40 


240 
170 
160 










Total 








43 -38 


1,292 


















98.43 


3,i97 











very limited variety, such as salt pork and "hard- tack," were 
liable to be afflicted with a peculiar and often fatal disease called 
scurvy. Afflicted persons rapidly recovered when supplied with 



Organic Compounds 



437 



fresh vegetables or even with the juice of oranges, lemons, or 
limes. The British navy and mercantile marine have for fifty 
years or more been required to provide sailors regularly with 
lime juice. Scurvy is now of rare occurrence. 

A still more remarkable case is found in the cause of and 
remedy for the peculiar oriental disease beri-beri. This affects 
people who live chiefly on a diet of rice and fish. Recent 
experiments, particularly with pigeons, has shown that these 
birds thrive on natural rice, while if fed on the polished grains 
they quickly suffer from the disease and soon die. Dangerously 
sick pigeons make marvelous recovery when given small amounts 
of the material removed from the grains in the process of polish- 
ing. Further experimentation has shown conclusively that 
beri-beri is the result of a diet deficient in a substance known 
as water soluble B found in the germ of the rice grain. This 
same substance is also found in the germ of wheat and other 
grains. 

Very extensive experiments have shown conclusively that 
rats cannot grow nor even live long on diets containing adequate 
amounts of proteins, fats, and carbo- 
hydrates if every article of food has 
been highly purified by chemical pro- 
cesses. This is because certain essential 
substances contained in the natural 
foods have been removed in the pro- 
cesses of purification. Milk is espe- 
cially rich in these essential substances. 
By adding small amounts of milk to 
the diet of rats living on purified food a 
wholly satisfactory ration is obtained. 
The results of such experiments are 
shown in Fig. 97. The lower curve 
shows the change in average weight of 
six young rats fed on a diet of purified foods alone. The upper 
curve shows the weights of the same number of young rats 
which received the same food ration as the first six but had in 
addition 2 c.c. each of milk per day. 



90 
70 

so 










30 







Days 

Fig. 97 



438 Introduction to General Chemistry 

In addition to water soluble B, milk contains also another 
peculiar substance called fat soluble A. The latter occurs also 
in the leafy parts of most vegetables. The term vitamines 
is usually used to designate the important substances, fat 
soluble A and water soluble B. It is now definitely established 
that animals, including man (especially children), must have 
for growth and health a constant supply of both these vitamines. 
A diet (for man) of roots, tubers, seeds, and meat may furnish 
sufficient proteins and have adequate fuel value and still be 
markedly deficient in vitamines, particularly fat soluble A. 
The deficiency is best avoided by the liberal use of milk and leafy 
vegetables; of these, milk is the more important. 

The chemical nature of the vitamines still remains to be 
discovered, but their importance as food constituents is no longer 
in question. There is little doubt that vitamines act cata- 
lytically. 

690. Dependence of Animals upon Plants. — All animals, 
including man, depend for their food upon plants or upon other 
animals which in turn feed upon plants. Animals cannot sub- 
sist upon the free elements that compose their food, nor even 
upon the simpler compounds of these elements, such as carbon 
dioxide, hydrocarbons (643), ammonia, amines (658), etc., but 
must have the far more elaborate compounds, the carbohydrates 
(682), fats (677), and proteins (685), and in addition mineral 
salts and vitamines (689). 

Plants, on the other hand, require for their sustenance far 
simpler materials, principally carbon dioxide, water, and simple 
nitrogen compounds, such as ammonia or nitrates, and also 
small amounts of mineral matter. Some of the simplest organ- 
isms contain species of bacteria, by aid of which they assimilate 
free nitrogen from the air (515). 

691. Photosynthesis. — The plant products formed from these 
simple inorganic substances have far more energy than the 
latter. What then is the source of this energy? Plainly the 
light and heat necessary for the growth of all plants (excepting 
fungi and other parasitic plants which live on decaying animal 
or vegetable matter). The energy of sunlight is transformed 



Organic Compounds 439 

in the growing plant into the chemical energy of the plant 
products. 

Cellulose and starch are the most abundant plant products. 
For the formation of these only carbon dioxide and water are 
theoretically required : 

6C0 2 + 5H 2 -> C 6 H I0 O s + 60 2 . 

However, this reaction does not take place in the simple fashion 
indicated by this equation. It takes place only in plants and 
then only in such parts as contain the characteristic green sub- 
stance chlorophyl. Probably several intermediate stages exist 
in the change of carbon dioxide into starch or cellulose; in any 
case oxygen is always a product of plant growth. Since this 
building up of complex products occurs by the aid of light the 
process is termed photosynthesis. 

692. Some Common Organic Explosives. — We have referred 
earlier (571) to the great practical importance of explosives. 
The most useful explosives are made by the nitration of organic 
substances such as glycerine (679), cotton (684), phenol (672), 
and toluene (668). The products thus obtained will now be 
briefly described. 

Nitroglycerine is formed when glycerine is added drop by 
drop to a cooled mixture of concentrated nitric and sulfuric 
acids. The nitroglycerine separates as an insoluble heavy oil 
when the product is poured into water. It is a violent explosive 
which finds extensive use in blasting. Nitroglycerine is the nitric 
acid ester (657) of the alcohol glycerine (679). Its formula is 

H 2 CON0 2 

I 
HCON0 2 

I 
H 2 CON0 2 

Its explosive nature arises from the fact that it contains more 
than sufficient oxygen to change all its hydrogen and carbon into 
water and carbon dioxide respectively. At the high tempera- 
ture reached in the explosion water is, of course, gaseous, as are 
likewise the other products, CO, and N>. Therefore the products 



44° Introduction to General Chemistry 

of the explosion occupy many times the volume of the original 
substance. This explains the enormous force produced by the 
explosion. Dynamite, which is a solid mixture of infusorial 
earth and nitroglycerine, is much safer to handle than the latter 
substance. 

693. Nitrocellulose, or guncotton, is made from cotton by 
a process similar to that used in making nitroglycerine. The 
product is a white solid scarcely differing in appearance from the 
cotton from which it is made. It is far less sensitive to shock 
than nitroglycerine, especially when in a moist state, and since 
it can be perfectly exploded while moist by the use of a detonator 
(573) it is one of the safest and most useful of explosives. Gun- 
cotton can be physically compounded with nitroglycerine to 
form cordite, a transparent solid resembling amber in appear- 
ance. This is one of the most important propellants for large 
projectiles. Modern smokeless powders are products closely 
related to guncotton. 

The complete nitration of phenol, C6H 5 OH (672), yields 
tri nitro phenol, or picric acid, C6H 2 (N0 3 ) 3 OH. This substance 
and its ammonium salt (673) are powerful explosives. Their 
most extensive use is in shrapnel. 

The explosive commonly designated as T.N.T. is tri nitro 
toluene. It has also been mentioned earlier (571). During 
the war it was used in enormous amounts in shrapnel and other 
explosive shells. 

A few words may be added at this point regarding the differ- 
ences in explosives, since the student will naturally wonder why 
so many different explosives are used. The most important 
properties which determine the character of an explosive are 
(1) its sensitiveness, (2) its force of explosion, determined by 
the volume and temperature of the gaseous products, and 
(3) its velocity of explosion. Of these three properties the last 
is very important; for although ordinary observation would 
indicate that every explosion is instantaneous, this is far from 
being the case. Every explosion requires time for its comple- 
tion. A satisfactory propellant must not explode too rapidly. 
It must allow time for the projectile to get under way, otherwise 



Organic Compounds 441 

it would burst the gun. For exploding shrapnel shells and for 
blasting rocks rapidly exploding substances are used. 

694. Other Products of Nitrocellulose. — When the nitration 
of cotton is not carried so far as in the preparation of guncotton 
the product is known as pyroxyline or soluble cotton. This 
substance resembles guncotton closely but is less explosive. 
It dissolves readily in many organic solvents, such as acetone 
(656) and esters (657), and in a mixture of alcohol (641) and 
ether (642). The solution so obtained is called collodion; it is 
used as an adhesive, as a liquid court-plaster, as a coating for 
incandescent gas mantles, and for many other purposes. 

Artificial silk, a product resembling silk in appearance, but 
totally different chemically, is made in one way from collodion. 
The process consists in " spinning" a concentrated collodion 
solution from a fine glass capillary by means of high pressure. 
Upon coming into the air the solvents evaporate at once, leaving 
a filament of the fineness of a natural silk fiber. These filaments 
are made into threads ; but this material is extremely combustible 
and must be denitrated by a chemical process which changes 
it back into cellulose without altering its beautiful silky luster. 
Artificial silk is made in large quantities. 

A rough way of telling the difference between artificial and 
natural silk is to set fire to a small piece of the fabric. Natural 
silk burns the way hair does, melting back of the flame and 
giving off a characteristic odor. Artificial silk burns as does a 
thin piece of cotton. It does not melt and gives virtually no odor. 

Celluloid is made by thoroughly kneading and rolling a warm 
mixture of pyroxyline and camphor, doHj 60. The latter is a 
product of the camphor tree. The many forms and uses of 
celluloid are too well known to require description. Photo- 
graphic films made from this material are highly inflammable. 
On this account celluloid is combined with less inflammable 
materials like cellulose acetate (acetic acid ester of cellulose) in 
the preparation of photographic films for moving pictures. 

Leather substitute, or artificial leather, is made by coating 
heavy cotton cloth with a preparation in which pyroxyline is 
the principal ingredient. 



442 Introduction to General Chemistry 

695. The Materials of Chemical Warfare. — The introduction 
by the Germans of chlorine in warfare was rapidly followed by 
the use of several other poisonous and irritating substances, 
practically all of which were carbon compounds. One of the 
most important of these was phosgene, or carbon oxychloride, 
COCl 2 . This gas is made by the union of carbon monoxide 
(632) and chlorine: 

CO+Cl 2 ->COCl 2 

It has a powerful, choking odor and is very poisonous. It is 
readily liquefied (boiling-point 8°). It was frequently mixed 
with liquid chlorine in gas attacks. 

The so-called mustard gas is not a gas at all but a colorless 
liquid known to chemists as dichlor diethyl sulfide (C1C 2 H 4 ) 2 S. 
It has a rather faint odor and at first appears harmless enough; 
but inhalation of the vapor or mist produced by explosion of a 
shell containing some of it causes terrific inflammation of the 
lungs, frequently proving fatal. Almost inconceivably minute 
amounts on the skin cause in the course of a few days deep and 
dangerous wounds. The substance is made from ethylene, C 2 H 4 , 
and sulfur chloride, S 2 C1 2 (601), and has the following structure: 

H H H H 

I I I I 

CI— C— C— S— C— C— CI 

I I I I 

H H H H 

It seems to act by being absorbed by the skin and then slowly 
hydrolyzing within the tissues to form hydrochloric acid and 
other products. 

Chloropicrin, CC1 3 N0 2 , made by the action of bleaching 
powder (351) on picric acid (673), is a high-boiling liquid. It 
is extremely irritating to the eyes, is poisonous, and causes 
vomiting. It passes through clothing and the fabric of a gas 
mask rather readily, and it is difficultly absorbed in the canister 
of chemicals used with a mask. This substance was used in 
large amounts toward the end of the war. Chloropicrin had 
no use before the war; it is now proposed to employ it as an 
insecticide. 



Organic Compounds 443 

696. Organic Synthesis.— The term organic synthesis means 
the artificial building up of organic compounds from simpler 
substances, ultimately from the elements. A century ago it was 
generally believed that the complex carbon-containing products 
of plants and animal organisms (organic compounds) were 
formed as the result of vital forces, and that it was impossible 
for chemists to make them artificially. In 1828 Wohler proved 
the fallacy of this idea by the synthesis of one of the most char- 
acteristic of animal products, urea, C0(NH 2 ) 2 (659). His 
method consisted in oxidizing potassium cyanide, KNC (665), 
to the cyanate KNCO by heating it with litharge, PbO : 

KNC+PbO->KNCO+Pb. 

From KNCO and (NH 4 ) 2 S0 4 he obtained by double decomposi- 
tion ammonium cyanate, NH 4 NCO. This salt when warmed 
in solution gradually changes into urea, 

NH 4 NCO->CO(NH 2 ) 2 . 

The announcement of Wohler's discovery created a profound 
sensation and led to the expectation that other vital products 
could also be synthesized. This expectation has in the past 
ninety years been realized far more completely than the most 
enthusiastic chemist of Wohler's day could have predicted; for 
not only have the greatest variety of vital products been synthe- 
sized, but thousands of related and new and unrelated organic 
substances have been made in the laboratory, so that today their 
number is legion. True, however, much remains to be done, 
for as we have seen in one instance the proteins have as yet 
not been synthesized. By way of illustrating the nature and 
scope of organic synthesis we shall in the following paragraphs 
recount briefly a few typical cases. 

697. Essential Oils and Perfumes. — Spices, fruits, and 
flowers owe their characteristic odors or perfumes to small 
amounts of substances which can in many cases be isolated and 
purified. These fragrant substances are volatile oils or solids 
known generally as essential oils. Frequently the oil from a 
given source is a mixture of several definite chemical substances, 
although in other cases but a single substance is present. 



444 Introduction to General Chemistry 

The first essential oil made synthetically was that contained 
in bitter almonds and well known as the oil of bitter almonds. 
Examination of the natural oil showed it to be benzaldehyde, a 
substance having the formula 

/\cb 

I I 
\/ 

When toluene (668) is treated with chlorine it gives benzal 
chloride, C 6 H 5 CHC1 2 . The latter, by treatment with water, 
gives benzaldehyde (670). The product so obtained has the 
same agreeable odor as that made from bitter almonds and is 
extensively used as a substitute for the natural oil in flavoring 
extracts. 

The chief constituent of oil of cinnamon, known as cinnamic 
aldehyde, has the formula 

C 6 H S CH = CHCH0. 

It can also be made synthetically from toluene, and when so 
prepared has the same fragrant odor as cinnamon. 

The characteristic constituent of vanilla is also an aldehyde, 
called vanillin, somewhat related to the two foregoing substances. 
Its formula is 

o 

CH 3 

It is a white crystalline solid, which is present in vanilla beans 
to the extent of about 1 per cent. After the discovery of its 
formula as just given it became possible to make the substance 
synthetically. It was soon found that other essential oils were 
closely related to vanillin. Thus oil of cloves consists largely of 



HO 



CH 2 CH = CH 2 



O 
CH, 



Organic Compounds 445 

This oil is easily and cheaply made from cloves. It can be 
converted into vanillin by a process of oxidation, which thus 
affords a good practical method of making the latter valuable 
substance. 

It is a well-known fact that organic substances with closely 
related formulae have similar properties, and this is beautifully 
illustrated by the near relatives of the two substances just con- 
sidered. The fragrant oil of sassafras consists largely of 
safrole, 

NCH 2 CH = CH 2 

Xo \/ 

This substance can be oxidized to an aldehyde piperonal, 

H 



H,C<¥' 



""OCT 



a substance having the delightful odor of heliotrope. Thus an 
exquisite and costly perfume is made in the laboratory from the 
cheap and abundant oil of sassafras. 

Many essential oils belong to the class of substances known 
as esters (657). For example, oil of wintergreen is the methyl 
ester of salicylic acid, readily made by following reaction: 

x )cOOH +C ^ OH -OcOOCH+^ 

The ester, which is made technically on a large scale, is a color- 
less oil having exactly the odor of wintergreen. It will be of 
interest to trace the manufacture of this substance from its 
beginning. To get salicylic acid we start with benzene (667), 
obtained from coal tar. This by treatment with sulfuric acid 
gives benzene sulfonic acid (672), 

|S0 2 OH 
/ 



446 Introduction to General Chemistry 

When the sodium salt of this acid is fused with an excess of 
sodium hydroxide it yields phenol or carbolic acid (672), 



r 



OH 
\/ 

The sodium salt of this substance reacts with carbon dioxide to 
form the salt of salicylic acid, 

C 6 H 5 ONa+ C0 2 -> C 6 H 4 OHCOONa, 

from which the free acid is readily obtained by the action of 
sulfuric acid. 

The subjects discussed in this section are of importance 
chiefly as a means of illustrating the ways in which some natural 
products are made in the laboratory. It is not worth while for 
the student to attempt to memorize the formulae or the reactions 
involved. The examples given are chosen from hundreds 
equally important and interesting. 

698. Dyes. The Synthesis of Indigo. — In earlier times 
dyes were usually plant or animal products. A few coloring 
substances, not dyes in the strict sense of the word, were of 
mineral origin. At present all but a small proportion of dyes 
are made artificially from substances obtained from coal tar. 
In a few cases dyes originally obtained from plants are now 
made in the laboratory. The most interesting case in point is 
that of indigo, which has long been one of the most important 
blue dyes. It is the product of a plant grown extensively in 
India, Java, and elsewhere. After the development of synthetic 
chemistry it was apparent that the artificial production of any 
natural chemical substance was possible, though doubtless in 
many cases difficult of realization. This possibility led chemists 
to hope that indigo could be made in the laboratory. The first 
step was to discover its structural formula. The purification 
of indigo proved easy, and the analysis of the pure crystalline 
substance was a matter of routine. The simplest formula 
possible according to the analysis was CgH 5 NO; but the vapor 
density (71, 217) indicated a molecular weight corresponding 
to twice this formula, namely, C l6 H I0 N 2 O 2 . Now it would be 



Organic Compounds 447 

possible to . think of hundreds of molecular arrangements of 
16 atoms of carbon, 10 atoms of hydrogen, and 2 each of nitrogen 
and oxygen, all in accord with accepted laws of valence (648); 
only experiment could decide which if any of these formulae 
represented the structure of indigo. Let us try to explain how 
the organic chemist attacks this kind of a problem, for he must 
solve it beyond doubt if he hopes to synthesize the substance. 
By distillation of indigo with KOH one obtains aniline (674), the 
structure of which is known to be 

/Nnh 2 

\/ 

The oxidation of indigo yielded a related substance, isatine, for 
which the structure had been shown to be 



( 



\/ n \ 



H 

\ 
C=0 



\/\c/ 

II 

o 

These facts indicate that the indigo molecule contains the atomic 
group 

,/\/ N \ 
J c= 

V\c/ 

II 

This group contains 8 atoms of carbon and 1 of nitrogen; since 
indigo is C I 6H I0 N 2 2 , the latter probably contains two isatine 
groups. Finally, after many years of research the structure of 
indigo was shown to be the following: 

TT TT 

/\/ N \ / N \/\ 

II c=c II 

\A C / \ c /\/ 

o o 



448 Introduction to General Chemistry 

After the structure of indigo had been found it was not long 
until several methods were devised by which it could be made 
synthetically; for once he knows its structure the organic 
chemist can attempt the building up of a molecule with nearly 
as great certainty of ultimate success as an architect can con- 
struct a building from prepared plans. A successful technical 
synthetic process must employ starting materials that are suffi- 
ciently abundant and cheap. For the artificial production of 
indigo we start with naphthalene (moth balls), C I0 Hg (669), the 
structure of which is 

\/\ 



f 



Oxidation of this substance yields phthalic acid, 

'Ncooh 



r 



wCOOH 

which in turn yields, by the aid of ammonia, a substance having 

the formula 

O 



I I NH 

V\c/ 



o 

o 



This by the action of sodium hypochlorite (350) gives an acid 
of the following composition: 



< 



\/ NUl 



\C00H 

The next step is the union of this acid with chloracetic acid, 

a /NH 2 /v /NHCH 2 COOH 

+ ClCH 2 COOH->| I +HC1 

^COOH \/\cOOH 






Organic Compounds 449 

When the product is heated with sodium hydroxide it is changed 
into jj 

/\/ N \ 

CH 2 

V\c/ 
o 

Upon oxidation with air, two molecules of this substance give 
water and a molecule of indigo, 

XT XT 

/\/ N \ / N \/\ 

II c=c II 

NAc/ \(/V 

o o 

Artificial indigo can now be made cheaper than the natural dye 
and has nearly driven the latter from the market. 

The story of indigo is closely paralleled by that of madder, 
an extract from the root of which dyes turkey red. The same 
dye is now made synthetically from anthracene, C I4 H I0 , a 
coal-tar product resembling naphthalene. 

The story of indigo has been told in some detail largely to 
illustrate the method by which a natural substance is reproduced 
synthetically. The steps in the process, which in all cases are 
the same, are the following: first, the isolation and purification 
of the substance; second, the analysis; third, the, molecular- 
weight determination to discover whether a multiple of the 
simplest possible formula is the true formula; fourth, the dis- 
covery of the structure; fifth and last, the synthesis from simple 
substances or from the elements. Of these steps the fourth is 
usually the most difficult, although both this step and the last 
demand the highest skill of the chemist. It would require the 
space of a chapter to give even a brief account of the coal-tar 
dyes. Suffice it to say here that about nine hundred different 
coal-tar dyes are made, and of these about three hundred 
are in active demand. Each is a definite chemical substance, 
the preparation of which is on the average as complicated a 
process as that of making indigo. At the beginning of the war 



450 Introduction to General Chemistry 

90 per cent of all dyes used in America were imported, largely 
from Germany. At the close of the war we were making in 
the United States a greater tonnage of dyes than we required, 
although many less important kinds of dyes were not yet made 
in this country. 

699. Synthetic Medicinals and Photographic Chemicals. — 
At the present time a large number of valuable medicinal sub- 
stances are produced synthetically from coal-tar products. 
Many of these substances are obtained from the same inter- 
mediate products, so-called intermediates,^ that are used in 
making dyes. Chemicals used as photographic" developers are 
also derived from similar intermediates. These facts give an 
added significance to the development of an American dye 
industry, since it means that we shall also be able to manufacture 
in this country our required medicinals and photochemicals. A 
long step in this direction was taken during the period of the war. 

700. Rubber. — It has long been known that the composition 
of natural rubber is represented by the very simple formula 
C 5 H8, but that the actual formula is probably C^H^ or perhaps 
(CioH^)^ Thus rubber has the same percentage composition 
as the comparatively cheap and abundant substance turpentine, 
CioH^ (248). It is therefore not surprising that enormous 
efforts have been made to produce synthetic rubber. On account 
of the great practical importance of this subject a brief discus- 
sion may prove of interest. One of the products of the distilla- 
tion of rubber is a colorless volatile liquid called isoprene. This 
has the simple formula C 5 H 8 and has been shown to have the 

structure 

CH 3 CH:CHCH:CH 2 . 

Isoprene has been made synthetically by several methods. It 
is a colorless liquid, which by treatment with certain catalytic 
agents changes into synthetic rubber: 

2 C 5 H8 -> C10H.16' 

As long ago as 191 2 a synthetic-rubber automobile tire that 
had run 5,000 miles and was still in good condition was exhibited 
at the International Congress of Applied Chemistry. 



Organic Compounds 451 

The practical problem of the synthesis of rubber is then 
resolved into the sufficiently cheap production of isoprene. The 
starting material must be a compound of carbon and hydrogen 
with or without other elements, which is cheap and abundant 
and convertible with a good yield into isoprene. No wholly 
satisfactory technical method has yet been found, for the simple 
reason that the cost of artificial rubber by any of the processes 
so far devised is greater than that of natural rubber. Here then 
is a most interesting and important problem for the future. To 
the chemist who solves it will come both fame and wealth. 

701. Sources of Organic Material. — The possibilities of 
synthetic organic chemistry are very alluring to one possessed 
of an active imagination; but it is necessary to keep in mind 
that all synthetic substances must be made from other carbon 
compounds, and that if a substance is to be made on a practical 
scale it can be made only from something sufficiently abundant 
and cheap. These facts make it worth while to consider the 
natural sources of organic materials. 

The principal sources are (1) plants, (2) animals, (3) coal, 
and (4) petroleum and natural gas. It is well known that coal 
is of vegetable origin; and it seems probable that petroleum 
comes from both vegetable and animal matter. Going back a 
step farther we recall that all animals are dependent on plants 
for nourishment, so that we find that plants are the ultimate 
source of all other organic material. We have also seen that plants 
obtain their carbon from the carbon dioxide of the air. Although 
the C0 2 content of the air is but three parts in 10,000 by volume, 
the weight of this gas above every square mile of the earth's 
surface amounts to about 10,000 tons. Therefore there is in the 
air an immense reserve of the element carbon. 

We have already referred to the many important constituents 
of coal tar (634, 667). These serve as the more immediate 
starting materials for numerous organic syntheses. The dis- 
tillation of wood also yields, in addition to charcoal, several 
very useful products, of which methyl or wood alcohol (645"), 
acetone (656), and acetic acid (653) are the most important. 



CHAPTER XXVII 

THEORY OF DILUTE SOLUTIONS 

Our purpose in this chapter is to develop in greater detail 
than before the picture of the condition of a dissolved sub- 
stance. This will be accomplished to best advantage by going 
over at the same time parallel work done on gases. 

702. The Reality of Molecules. — According to the kinetic 
theory gases are composed of tiny particles moving at random 
with velocities of rifle bullets and spaced from each other dis- 
tances which on the average are large compared with the diameter 
of the particles (196). This picture has given us an understand- 
ing of why all gases diffuse through and occupy any space in 
which they are released. It has also explained why all gases 
follow the same laws, namely those of Charles and Boyle 
(4, 5). Further than this, however, the kinetic theory has 
explained why under great pressures and low temperature gases 
do not follow these laws closely (225). So great has been 
the assistance given by the theory in understanding the 
behavior of gases that the actual existence of molecules 
has seemed to many beyond the possibility of doubt, even 
though the molecules themselves have never been seen. More 
recently, however, the crowning evidence in proof of the 
reality of molecules has been brought out by work with the 
ultra-microscope. 

703. Perception of Molecules. — The celebrated French 
physicist Perrin, to some of whose work on the cathode rays 
we have referred earlier (478), says, " Direct perception of the 
molecules in agitation is not possible for the same reason that 
the motion of the waves is not noticed by an observer at too 
great distance from them. But if a ship comes in sight, he is 
able to see that it is rocking," and this "will enable him to infer 
the existence of a possibly unsuspected motion of. the sea's 
surface." While we cannot see molecules themselves, we can 

452 






Theory of Dilute Solutions 



453 



by devices now to be described see the phenomenon which is 
the counterpart of the boat rocking on the invisible waves. 

In place of a ship buffeted about by waves we shall observe 
particles of smoke floating in the air and exposed to the bombard- 
ment of air molecules. Although these smoke particles are 
far larger than the largest molecules (each consisting of many 
molecules), they are still too small to be seen even under a 
microscope unless they are brilliantly illuminated. The smoke 
(from a cigarette or the like) is contained in a small glass box 
beneath the objective of the microscope and is illuminated by a 
beam of light from a projection lantern focused at a point in 




r^gMMBBCS 



F 



A 



Fig. 98 



the smoke just under the objective. The arrangement is illus- 
trated in Fig. 98. The observer looks down through the micro- 
scope upon tiny smoke particles which are illuminated against 
the dark background by the horizontal rays from the lantern. 
The field of the microscope seems covered with a "milky way" 
of bright stars which are the flashes of light reflected to the 
observer's eye by the particles. Thanks to the dark background 
these can be easily distinguished. 

The flat-sided jar of water, N, shown in the figure is inter- 
posed in the beam from the lantern in order to cool it as much 
as possible before allowing it to strike the box; but even so con- 
vection currents are set up within the box. As a consequence 
of these currents all the particles drift along in one direction, 
but if the observer watches a single particle it will be seen to be 
dancing about in a most erratic way. This jiggling motion 
keeps up endlessly, as long as the particles themselves arc in 



454 Introduction to General Chemistry 

suspension. The smaller the particles the more violent is the 
motion. In more refined apparatus still smaller particles can 
be seen; these dart over distances many hundreds of times their 
own diameter. The fact that the motion is absolutely erratic, 
and that any suspension of any substance, the individual particles 
of which are very small, exhibits this same motion, shows that 
the latter is due to the irregular bombardment of the particles 
by the molecules of the gases in which they are suspended. The 
particles must be small enough to be moved appreciably by a 
few more hits by molecules on one side than on the other. This 
motion, called from its discoverer Brown the Brownian move- 
ment, is the phenomenon which is the counterpart of the rocking 
of the boat upon the invisible waves. 

704. The Kinetic Theory of the Liquid State. — In comparing 
liquids with gases we note at once the far greater density of the 
former. The kinetic theory (198) describes molecules in the 
liquid state as moving about, much as they do in the gaseous 
state, but as greatly hampered in their motion on account of the 
crowding of the molecules. In liquids the average path between 
collisions is much shorter than in gases, so that diffusion through 
the former is very much slower than through the latter. The 
smaller distance between the molecules means that the attrac- 
tion between them is of greater moment than in the gaseous 
state. This attraction falls off rapidly with increasing distance, 
so that each molecule is affected by its immediate neighbors 
only. As a result, within the liquid this effect is balanced, since 
the pull exerted by the neighbors comes equally from all sides. 
Only at the surface of the liquid does this force become one- 
sided and so hampers the free motion of the molecules. As a 
result only a few of the •faster-moving molecules can escape 
from the surface (199). According to the kinetic theory the 
average kinetic energy of molecules in solution is the same as 
that of molecules in the gaseous state at the same temperature, 
so that the light molecules in solution are to be pictured as mov- 
ing very much faster than the heavy ones. In many ways the 
success of the kinetic theory in helping us to understand the 
liquid state has been so marked that its validity was accepted 



Theory of Dilute Solutions • 455 

even before the development of the more direct evidence of the 
ultra-microscope. 

705. Brownian Movements in Liquids. — The existence of 
molecular motion in liquids can be demonstrated by the Brownian 
movements of very tiny particles in suspension in liquids, just 
as it can be in the case of smoke particles. Suspended particles 
of suitable size, whether of clay, lampblack, starch, or anything 
else, all exhibit these movements irrespective of their composition. 

The pigment gamboge, when rubbed in a little water or 
dissolved in alcohol and added to water, gives a yellowish turbid 
liquid which consists of suspended droplets of suitable size to 
show the Brownian movements. A microscope with oil-immersion 
objective is best used. If a drop of the gamboge suspension is 
mounted for observation under the objective and viewed with 
the ordinary illumination reflected from the mirror below the 
stage of the instrument many tiny particles will be seen in the 
field. Careful observation will show that the smaller particles 
are in a constant tremor, moving erratically over paths approxi- 
mately one to several times their own diameters in length. It 
seems difficult for the observer to believe that these moving 
particles are not alive, so lifelike is their motion! 

706. The Revelations of the Ultra-Microscope. — Suspensions 
of far smaller particles may be observed if instead of the ordinary 
microscope we use the ultra-microscope. To build a crude form 
of the latter we may employ the same set-up as was used for 
the examination of the smoke particles, except that we must 
use an oil-immersion lens and a so-called condenser. In this 
case the light from the lantern is again made to pass through 
water, but this time, instead of being directed next against the 
object, it is made to strike the mirror below the stage of the 
microscope and is reflected into the condenser, which directs it 
obliquely against the droplet on the slide, as shown by the path 
of the arrows in Fig. 99. A drop of pure water placed on the 
slide appears only as a dark field to the observer; but if there 
are tiny particles present in it flashes of light are reflected from 
them to the eye. Zsigmondy, who with Siedentopf perfected 
the ultra-microscope, describes his first view of a fine suspension 




456 Introduction to General Chemistry 

of gold particles in water thus: "The small gold particles move 
and that with astonishing rapidity. A swarm of gnats in a 
sunbeam will give one an idea of the motion of the gold particles. 
.... They hop, dance, jump, dash together and fly away from 
each other so that it is difficult to get one's bearings." The very 
fine particles seen under the ultra-microscope are observed to 
travel much more rapidly and to cover 
greater distances than those previously de- 
scribed. It should be realized that any 
suspension in any liquid shows this same 
erratic motion. Specimens have been under 
observation for years, but no abatement 
of the Brownian movement has been ob- 
served. Suspensions in liquids taken from 
Fig. 99 the interior of quartz crystals show these 

same motions, and yet they are thousands 
of years old 1 We have, therefore, the same type of evidence of 
the reality of the existence of swift-moving molecules in liquids 
as we have in gases. 

707. Visible Atmospheres. — Perrin was able to show that 
fine suspensions may be considered as visible atmospheres. The 
pressure due to the impacts of the ultra-microscopically visible 
particles is proportional to the concentration and to the absolute 
temperature, just as the pressure due to gas molecules is pro- 
portional to their concentration and to the absolute temperature. 
The actual number (194) of molecules in a gram molecular 
weight of a substance in the gaseous state has been estimated 
as 6X 10 23 . Perrin estimated the number of suspension particles 
in what would be a gram molecular weight of a suspended sub- 
stance, namely the weight of substance in 22.4 liters of a suspen- 
sion the particles of which at o° exert a pressure of 760 mm. 
The numbers found for a great variety of suspensions were 
close to 6 X 10 23 . The viscosity of the suspension medium seemed 
to make no difference in the result. It did not matter whether 
the suspension medium was pure water or the very viscous 
glycerine. The average weight of the particles was varied 
70,000 fold, and yet the results were the same. Thus the gas 



Theory of Dilute Solutions 



457 



laws and Arogadro's hypothesis apply to these suspensions. 
We are then to think of fine suspensions as miniature atmospheres 
of colossal molecules which are actually visible. 

708. Particles in Solution vs. Particles in Suspension. — 
Particles visible in the ultra-molecular microscope are called 
suspensions, but particles just too small to be seen in this way 
must be counted as in solution. The change from solution to 
suspension should be thought of as gradual rather than abrupt, 
for the distinction seems to be fundamentally one of the size of the 
particle. 

709. Substances in Solution vs. Gases. — The knowledge that 
particles in suspension in a liquid act like huge gas molecules 
inclosed in the same volume as that of the solution leads at once 
to the conception that substances in solution may also be con- 
sidered in the same light. To test this theory the pressure of 
the dissolved molecules must be measured separately from that 
of the molecules of the solvent. This is obviously not a simple 
undertaking, since the full impacts of molecules in solution are 
not felt by a surface in contact with the latter 
because of the attraction inward of the molecules 
which come to the surface layer. The same force 
which prevents the free migration of molecules 
from the surface of the liquid prevents the 
measurement of the force of their impacts by such 
devices as are used to measure gas pressure. 
However, devices to suit the purpose have been 
found, and these will now be studied. Since 
they depend on the operation of a semipermeable 
membrane, a simple example of the latter will first 
be described. 

710. Semipermeable Membranes. — A so- 
called semipermeable membrane is a membrane 
which is permeable to one of two substances in 
a mixture but not to the other. A layer of 

thin cloth (muslin) stretched over the tubulated end of the oil 
manometer, as shown in Fig. 100, and wet with water will serve 
as a membrane permeable to ammonia but not to air. The 



NH 3 



Fig. 100 



45! 



Introduction to General Chemistry 



stopper at A permits the insertion of red litmus paper into the 
inner chamber, as shown in the figure. When ammonia is col- 
lected in a beaker inverted over the membrane, the oil in the 
manometer moves, showing an increase of pressure in the inner 
chamber. Meantime the litmus paper is seen to turn blue, thus 
showing that ammonia has entered. This it has done by dis- 
solving in the upper surface of the water film and diffusing 
through it. As a result some ammonia molecules soon appear 
at the inner surface of the film and then 
escape into the air within the apparatus, just 
as in the case of the diffusion through a gas 
(191) the migration of ammonia molecules 
through the water and air layers proceeds 
spontaneously from regions of higher concen- 
tration to those of lower concentration until 
there is no longer any difference between the 
two. Meantime the air inclosed in the ap- 
paratus is not able to escape, since it/is not 
appreciably soluble in water, and therefore 
the total pressure increases. 

711. The Solution Cell. — A similar appa- 
ratus for use with solutions may be easily 
prepared for demonstration purposes by at- 
taching a parchment-paper thimble to a glass 
tube, as shown in Fig. 101. The thimble is 
filled with concentrated sugar solution colored 
with cochineal and immersed in a beaker of 
pure water. Water can pass through the 
parchment wall in either direction, but the 
sugar and the cochineal cannot do so appre- 
3 ciably. Since the color can only accompany 
Fig. ioi the sugar, it serves as a convenient marker 

for the presence of the latter. Soon after 
the cell is filled and the stopcock of the filling tube, A, is 
closed the level of the sugar solution begins to rise in the verti- 
cal tube. In a few hours the surface of the sugar solution will 
be three or four feet above the level of the water in the outside 



Theory of Dilute Solutions 



459 



beaker. There is obviously a close analogy between this experi- 
ment and the preceding one. The water in the former cor- 
responds to the ammonia in the latter, since each can pass 
through .the membrane. The sugar, like the air, cannot pass 
through the membrane. Whatever the mechanism by which 
the cell operates, it is plain that the weight of the column of 
solution is borne by the membrane, except for the slight assist- 
ance given by the pressure exerted against it by the water in the 
outside beaker. If the column is allowed to get too high the 
thimble will be broken. For accurate measurements stronger 
membranes must be used, and they" must also be of substances 
more truly impermeable to sugar. The best membranes have been 
made by precipitating copper ferrocyanide 
in the pores of unglazed porcelain cups. 

A cell of this type is shown in Fig. 102. 
The porcelain cup which holds the sugar 
solution is fitted with a mercury manom- 
eter, the closed end of which contains nitro- 
gen gas. As water enters the cell through 
the copper ferrocyanide membrane deposited 
in the pores of the cup, the nitrogen in the 
manometer is compressed. Since the manom- 
eter is very narrow, only a small increase of 
volume on the part of the solution in the cell 
results in a very great compression of the 
nitrogen. The change in volume of the latter 
(together with the very small difference in 
levels of the mercury in the manometer arms) 
gives the necessary data for calculating the 
pressure within the cup. When no further 
change in volume of the nitrogen is seen, the pressure of the 
latter is just sufficient to check the inflow of water through the 
membrane. The pouring of the water into the sugar solution 
through the membrane is called osmosis, and the pressure which 
checks this flow is called the osmotic pressure of the solution. 

712. Results of Osmotic-Pressure Measurement. — Table 
XXXIII gives the osmotic pressures of cane-sugar solutions 




Fig. 102 



460 



Introduction to General Chemistry 



at 6o° at the concentrations indicated. From indirect evidence 
chemists long ago made up their minds that the molecular 
weight of sugar is 342, corresponding to the formula Ci 2 H 22 On 
(682). Hence we may give the concentration in percentage by 
weight of sugar in solution and also in terms of gram molecular 
weights of sugar to 1,000 grams of water, i.e., weight-molar con- 
centration. The osmotic pressure is given first in lbs. per square 
inch, in order that the magnitude of these pressures may be 
visualized, and then in terms of atmospheres. From the table 
it is evident that even for a 3 per cent solution the membrane 



TABLE XXXIII 

Osmotic Pressures of Sugar Solutions 



Concentration 


Osmotic Pressure 


Gas Pressure 

in 
Atmospheres 


Ratio 


Percentage 


Weight-Molar 


Lbs. Per Sq. In. 


Atmospheres 


O.P. to G.P. 


3-3 


O. I 


40 


2 . 72 


2 . 72 


1 .00 


6.4 




2 


80 


5-44 


5-43 


1 .00 


9-3 




3 


120 


8.14 


8.15 


1 .00 


12 .0 




4 


160 


10.87 


10.87 


1 .00 


14.6 




5 


202 


13.67 


I3-58 


1 .01 


17. 1 




6 


244 


16.54 


16.30 


1 .01 


19-3 




7 


285 


19.40 


19.02 


1.02 


21-5 


O 


8 


329 


22.33 


21.73 


1.03 



must stand a pressure of 40 lb. to the square inch. For a 
20 per cent solution the membrane must stand a pressure of 
over 300 lb. per square inch. Students will readily appreciate 
the experimental skill necessary to make these measurements, 
when they think of their own difficulties in securing an apparatus 
tight enough to stand less than a pound of pressure without 
leaking ! In the fifth column is given the pressure which a like 
molecular concentration of hydrogen' would exert at the same 
temperature. The close agreement between osmotic pressure 
and gas pressure is at once apparent from a comparison of the 
fourth and fifth columns. This is brought out in the sixth 
column, which represents the ratio of osmotic pressure to gas 
pressure. Where the ratio is equal to 1.00, the two may be 
regarded as equal. As has been said, this data is for solutions 



Theory of Dilute Solutions 461 

at 6o°. How osmotic pressure compares with gas pressure at 
other temperatures is shown in Table XXXIV, which represents 
the ratio of osmotic pressure to gas pressure for solutions at 
other temperatures as well. The data for both tables have been 
adapted from the very fine work of Morse and Frazer, of Johns 
Hopkins University. 

In both Table XXXIII and Table XXXIV it is apparent 
that the osmotic pressures are indeed close to the corresponding 
gas pressures. If we look at the deviations from the gas laws 
by noting how much greater than 1 . 00 the ratios are, we find 

TABLE XXXIV 

Ratio of Osmotic Pressure to Gas Pressure at Various Temperatures 



Concentration 



20" 



40 



6o° 



0.1 (1. 11) 

0.3 1 . 06 

0.5 1.07 

0.8 1 . 09 

I.O I . 12 



I.08 

I.06 
I .07 
I .09 
I .12 



I .OO 
I .02 

I -05 
I .07 
I .09 



I .OO 
I .OO 
I .OI 
I.03 
I.04 



that they are larger the lower the temperature and the more 
concentrated the solution. This reminds one of the deviations 
of the behavior of the gases themselves from strict agreement 
with the gas laws (225). 

713. Interpretation of Osmotic-Pressure Measurements.— 
How are we to explain the striking results of osmotic-pressure 
measurements? The fact that the measured osmotic pressure 
of a solution equals the pressure due to a gas at like temperature 
and of like molecular concentration as that of the dissolved 
substance in question is a coincidence so remarkable that we 
cannot call the agreement one of chance. i\pparently the cell 
operates because water is at a different concentration on the 
two sides of the membrane just as the air-ammonia cell operates 
because ammonia which could pass through the water mem- 
brane was at first more concentrated on one side than on the 
other. In the sugar-water cell presumably the pressure due 
to the impacts of the molecules against the membrane is at first 



462 Introduction to General Chemistry 

the same from both sides; hence the pressure of water molecules 
inside the cell is short of that outside by just the pressure due 
to the impacts of the sugar molecules. This pressure must be 
put upon the solution in order that the water may be made 
to pass back from solution to pure water at the same rate that 
it is coming into the cell and equilibrium thus be established. 
In this way we may explain the operation of the cell and the 
peculiar values found. 

In our previous work we have always given concentrations 
of solutions in terms of weight of dissolved substance per liter of 
solution. In the above-mentioned work concentration is given 
in terms of weight of dissolved substance per liter of solvent, 
since the values so obtained are a little nearer to the correspond- 
ing gas pressures, presumably because the amount of solvent 
is a measure of the actual free space through which the dissolved 
particles move. 

In measuring the pressure which must be put upon a solu- 
tion to check the inflow of solvent into it when separated from 
the solvent by a semipermeable membrane, we are measuring 
a force which is equal to the pressure exerted by the impacts 
of the molecules of the dissolved substances. From such 
measurements we learn that substances in dilute solution act 
much as though they were in the gaseous state, at the same 
temperature and inclosed in volumes equal to the volume of the 
pure solvent in the solution. 

714. Other Theories of Osmosis. — The theory just presented 
of the mechanism by which the semipermeable membrane cell 
operates is not the only possible explanation of the facts. For 
instance, some prefer to regard the motion of the solvent through 
the membrane as due to an attraction by the solution. Osmotic 
pressure would be the result of a sort of suction. As yet the 
theory presented above has been by far the more useful. 

715. Van't Hoff's Theory of Solutions.— The great Dutch 
chemist Van't Hoff was the first to point out that the gas laws' 
should hold for substances in very dilute solution. He drew 
his conclusions from considerations of a purely theoretical 
nature at a time (1887) when the data at hand for the experi- 



Theory of Dilute Solutions 463 

mental confirmation of his theory were far less accurate than 
those given in Tables XXXIII and XXXIV and showed far less 
clearly the truth of his statements. Van't Hoff pointed out 
that the laws of gases should be followed by substances in very 
dilute solution only, since in the more concentrated solutions 
the formation of complexes between the solvent and the dis- 
solved substance and between the molecules of the latter make 
conditions which are not parallel to those found in gases which 
follow the simple gas laws. 

716. The Avogadro-Van't Hoff Hypothesis. — Van't Hoff also 
asserted the so-called Avogadro- Van't Hoff hypothesis, a generali- 
zation of tremendous importance in the growth of modern chem- 
istry. This hypothesis is simply the application of Avogadro's 
hypothesis to substances in solution. Equal volumes of solu- 
tions of all substances which have equal osmotic pressures at the 
same temperature contain the same number of dissolved molecules. 

717. Molecular Weights from Osmotic-Pressure Data. — 
As a first result of the establishment of this principle we have 
learned how to find the relative weights of molecules in solution 
(molecular weights) by rinding what weight of the substance in 
question will exert 76 cm. of osmotic pressure at o° when dis- 
solved in 22.4 liters of solvent. 

718. Raoult's Work. Shorter Methods of Determining 
Molecular Weights. — Before Van't Hoff's work was published, 
Raoult, a French chemist, working with a great variety of sub- 
stances the molecular weights of which were known, had shown 
that a gram molecular weight of each dissolved in a given weight 
of solvent produced about the same lowering of the vapor 
pressure of the latter, provided the dissolved substance was not 
volatile. Van't Hoff pointed out (on the basis of reasoning 
which we need not now consider) that this is a necessary con- 
sequence of the fact that Avogadro's hypothesis applied to dilute 
solutions. Therefore the molecular weight of an unknown 
substance in solution may be determined by finding what weight 
of this substance is required to give a solution of .vapor pressure 
equal to that of a half-molar solution of a substance the molec- 
cular weight of which is known. Twice the value found would 



464 Introduction to General Chemistry 

be the[molecular weight of the new substance. Raoult's observa- 
tions are for dilute solutions, so that the work should not be 
done' with concentrations greater than half-molar. " 

We have already seen that one result of the lowering of the 
vapor pressure of a liquid by dissolving a non-volatile foreign 
substance in it is that the solution boils at a higher temperature 
than the pure solvent (128). Raoult showed that the elevation 
of the boiling-point is proportional to the number of gram molec- 
ular weights of substance dissolved in unit weight of solvent. If 
one gram molecular weight of a non- volatile substance is dis- 
solved in 1,000 g. of water the boiling-point is raised 0.50 . 
This is the so-called molar elevation for water. A gram molec- 
ular weight of a substance dissolved in 1,000 g. of ether elevates 
the boiling-point of the latter 2 . 20 . Here then is another way 
of determining molecular weights of substances in solution. We 
have only to find how much of the new substance must be dis- 
solved in 1,000 g. of solvent to produce the molar elevation of 
the boiling-point, in order to find the molecular weight of the 
substance in question. 

The freezing-point of a solution affords still another method. 
The depression of the freezing-point of the solvent, like the elevation 
of the boiling-point, is proportional to the number of gram molec- 
ular weights present per unit weight of solvent, provided the latter 
is the only substance which separates from solution during the 
freezing. Water solutions which contain one gram molecular 
weight of dissolved substance per 1,000 g. of water freeze at 
— 1. 8 5 . Hence the molar depression of the freezing-point 
of water may be taken as 1 .85°. That of the solvent benzene 
(667) is 5. 3 . Accordingly we need only find what weight of 
substance is necessary to produce the molar depression of the 
freezing-point of a suitable solvent to find the molecular weight 
of the substance in question in the solvent used. 

719. The Importance of the Work of Van't Hoff and of 
Raoult. — The general acceptance of Avogadro's hypothesis 
(about i860) initiated a new era in chemistry. A second great 
development of chemistry began in 1887, with the announcement 
of the Avogadro-Van't Hoff hypothesis. The latter furnished 



Theory of Dilute Solutions 465 

the basis for the understanding of reactions in dilute solution, 
just as the original hypothesis had done for gas reactions. Much 
of the rapidity of the development of the work on dilute solutions 
is due to the relatively simple experimental methods of Raoult 
for determining molecular weights of dissolved substances. The 
immediate result of the work of these two men was the presenta- 
tion of new and independent evidence concerning the ionic hy- 
pothesis which was just being developed by Arrhenius (in 1887), 
mainly from work on conductivity measurements. 

720. Molecular Weights of Acids, Bases, and Salts. — When 
the molecular weights of acids, bases, and salts were determined 
in water solution abnormal results were found. Thus in the 
case of hydrogen chloride, which as a gas has the molecular 
weight 36.5, the molecular weight in water solution was found 
to be much less than this, and furthermore, when measurements 
were made on solutions of greater and greater dilution, the 
results found grew smaller, reaching about 18.3 as a limit. 

The significance of abnormally low molecular weight in the 
case of ammonium chloride vapor (529) was shown to be due 
to dissociation into ammonia and hydrogen chloride. Arrhenius 
was quick to point out that the abnormally low molecular 
weights of electrolytes must also be due to the fact that the 
original molecules must have dissociated into ions in water 
solution. The" fact that these molecular weights get smaller 
with increasing dilution is evidence of increased dissociation 
with dilution; but finally, when the degrees of dissociation 
were calculated from freezing-point data and compared with 
the values found for the same concentrations by the conductivity 
method (408), the two were found to agree so closely that little 
doubt remained of the correctness of the ionic hypothesis. 

721. The Degree of Dissociation from Molecular- Weight 
Determinations. — The method of calculation of the degree of 
dissociation from molecular-weight determinations is best illus- 
trated by an example. In the case of potassium chloride the 
simplest formula is KC1, indicating a molecular weight of 74.6. 
If KC1 ionizes thus, 

KC1^K++C1-, 



466 Introduction to General Chemistry 

it is plain that if 100 molecules were originally present and 
80 per cent of them broke down into ions, the total number of 
free particles present would be 180 instead of 100. If we start 
with 74.6 g. of KC1 in 1,000 g. of water and 80 per cent is 
dissociated, we shall have present 0.2 gram molecular weight 
of KC1, 0.8 gram molecular weight of potassium ion, and 
0.8 gram molecular weight of chlorine ion, making a total of 
1 . 8 gram molecular weights from the 74 . 6 g. of salt dissolved 
instead of one gram molecular weight. The observed depres- 
sion of the freezing-point of water will be 1 . 8 times as great as 
is expected for one gram molecular weight of total substance 
dissolved in the same volume. Hence the apparent molecular 
weight of KC1 calculated from the freezing-point of a solution 
80 per cent ionized will be 74.6-5-1.8 = 41.4. Conversely, if 
the molecular weight of a KC1 solution found in the usual way 
by the freezing-point method is 41 .4, the degree of dissociation 
of the solution is found thus: 74.6-5-41.4 = 1.80. From this 
we conclude that the dissociation is 80 per cent. 

722. Degree of Dissociation of Electrolytes by Two 
Methods. — The degrees of dissociation of a number of electro- 
lytes at different concentrations as calculated from molecular- 
weight determinations by the method just illustrated are shown 
in Table XXXV (line I for each substance) in comparison with 
like data for the same substance as found from electrical con- 
ductivities (line II). (In the more concentrated solutions the 
latter data have been corrected for changes in the viscosity of 
the solution with concentration. Obviously, if the viscosity of 
the solution increases, the rapidity with which the same ions will 
move under the same attractive force will be diminished so that 
a correction is necessary.) Concentrations are given in terms of 
molecular weights per liter except in the cases of salts having 
bivalent ions. In these cases the fraction J indicates that the 
concentration is one-half the concentration indicated at the head 
of the column. The solutions represented in the same column 
contain equivalent weights of materials (403). 

Apparently the agreement between the results of the two 
methods is within a few per cent for concentrations up to 



Theory of Dilute Solutions 



467 



0.5 molar for salts consisting of two univalent ions in com- 
bination, and for concentrations up to 0.25 molar for salts 
consisting of two univalent ions united to one bivalent ion. 
With increasing concentrations for both types the divergences 
become larger. Data for salts made up of higher valent ions 
show even greater disagreements in the two methods. However, 



TABLE XXXV 

Degrees of Dissociation* 



Concentrations 


O.OI 


0.05 


O.IO 


0. 20 


0.50 


KC1I 

KC1II 


94% 
94 

94 
94 


89% 
89 

89 
88 

90 
89 

93 
94 

9i 
94 

82 
80 

84 
80 

72 
7i 

78 
77 


86% 
86 

88 
85 

86 
86 


83% 
83 

83 
82 

84 
83 


80% 
78 


NaClI 


82 


NaCl II 


77 


KBr I 


81 


KBr II 


94 

97 
97. 

96 
97 

88 
88 


77 


HC1I 




HC1 II . . . 








HNO3 I 








HNO3 II 








BaCl 2 (|) I 


79 
76 

82 
76 

65 
64 

73 

72 


76 

72 

80 
73 

57 
56 

67 
67 




BaCLQ) II 




CaCl 2 (|) I 




CaCl 2 (£) II 






Pb(N0 3 ) a (i)I 

Pb(N0 3 ) a <*) II 

K 2 S0 4 (^) I 


85 
85 

90 
87 


43 
45 

^7 


K 2 S0 4 Q) II 


62 







* Adapted from the work of Noyes and Falk. 

the complications which arise under both these circumstances 
make these discrepancies appear to be in keeping with the funda- 
mental theory. 

723. Summary. — The material developed in this chapter 
may be summarized as follows: The existence of free-moving 
molecules in liquids as described by the kinetic theory has boon 
put beyond the possibility of doubt by the discover)' that all 
very fine suspensions are in constant motion. Perrin has shown 



468 Introduction to General Chemistry 

that these fine suspensions behave like atmospheres of huge 
molecules. The kinetic energy of the suspension particles is 
equal to that of gas molecules at the same temperature. In fact, 
the gas laws may be applied to these suspensions. Even when 
the average mass of the molecules is varied 70,000 fold, the 
same results -are found. Since the distinction between an 
extremely fine suspension and a solution is fundamentally one 
of the size of the particles, it seems plausible that a substance in 
dilute solution might also be found to behave as though in the 
gaseous state. 

The pressure exerted by the molecules of a dissolved substance 
is determined by measuring an equal and opposite pressure, 
namely that of the inflow of solvent into the solution through 
a membrane permeable to the solvent but not to the dissolved 
substance. The pressures so measured are found to be very 
nearly equal to the pressure which a gas would exert if of the 
same molecular concentration and at the same temperature as 
the dissolved substance. Thus the analogy between substances 
in dilute solution and in the gaseous state is completely borne 
out. Van't Hoff, who was the first to develop the theory of 
dilute solutions, pointed out that Avogadro's Law should hold 
for substances in dilute solution. On the basis of this assump- 
tion the molecular weight of a substance in solution can be 
obtained by finding what weight of the latter dissolved in 22 .4 
liters will exert 760 mm. osmotic pressure at o°. Raoult, work- 
ing with a great variety of substances, the molecular weights of 
which were known, showed that a gram molecular weight of 
each produced about the same change in the vapor pressure of 
the solvent. The corresponding changes in the boiling-point 
and freezing-point can be easily measured. The theoretical 
deductions of Van't Hoff showed that the experimental methods 
of Raoult were simply other methods for determining the rela- 
tive weights of the molecules as they existed in the solutions. 

On the basis of the work of these two men the existence of 
particles smaller in weight than the original molecules of elec- 
trolyte dissolved in water solutions was established. Further- 
more the amounts of these smaller particles were found to be 



_ 



Theory of Dilute Solutions 469 

in agreement with the requirements of the ionic theory. . As a 
consequence the discovery of these facts put the ionic theory 
on a firm foundation. 

724. Osmosis in Nature.— Osmosis is one of the important 
processes operative in nature. This is illustrated by the effect 
of water on wilted plants. The latter have lost water in the 
air, but upon being put into water they "freshen" because the 
water pours through the cell walls and thus puts a pressure 
against the latter which stiffens them. This pressure could not 
accumulate if the cell walls were not only membranes, permeable 
to water, but also impermeable to much of the other contents 
of the cell. If a fresh flower is put into concentrated sugar 
solution, in a very few minutes it will be found to be badly 
wilted. This occurs because the concentration of water is less 
in the sugar solution than it is in the plant cells (713). The 
effect of sugar in withdrawing the juice from fruit is another well- 
known phenomenon of the same type. The foregoing are easily 
observed examples. However, many less easily observed but 
important physiological processes both of plants and animals 
depend on osmosis. 



CHAPTER XXVIII 
DISPERSE SYSTEMS 

725. Introduction. — The surface layer of a massive solid or 
of a liquid, as it is contained in an ordinary vessel, is of negligible 
importance chemically compared to its total mass. But if the 
solid is made into dust or the liquid is transformed into fine 
droplets, the surface layer is enormously increased and new 
phenomena become prominent. Dusts, mists, and other finely 
divided materials may be characterized as disperse systems. 
They are to be the subject of our next study. 

726. Cohesion and Adhesion. — The structure of gases and 
solids need not be discussed further here. But some points about 
the structure and behavior of liquids may well be taken up in 
preparation for the subsequent discussion. We are familiar 
with the sight of a liquid falling in drops from the tip of a pipette 
or burette. As soon as a given volume of liquid is free from 
outside influence, such as gravity, attraction from the walls 
of the containing vessel, etc., it tends to -form a sphere. If the 
drop is very small, the sphere will be nearly perfect even when 
pulled down against a supporting surface by gravity. This is 
well illustrated in the case of small drops of mercury resting on 
paper. The attraction between molecules falls off very rapidly 
as the distance between them increases. As a consequence 
only the immediate neighbors of a given molecule in a liquid 
exert any appreciable influence upon it, and, except in the surface 
layer, this attraction is balanced, since it comes equally from 
all sides. But in the surface layer the molecules are pulled 
inward and there is no compensating pull outward from beyond 
the surface. The net result is that the surface layer acts as a 
compressing membrane upon the rest of the drop. Work must 
be done against this inward pressure if a liquid is to be sub- 
divided. This fact may be easily observed if the finger is used 
to divide a large drop of mercury resting on a flat glass surface. 

470 



Disperse Systems 471 

The floating of a needle on the surface of water is another well- 
known experiment, used to illustrate this resistance of a liquid 
to deformation. When two drops touch each other they 
promptly merge into a single drop. This occurs because at 
the instant of contact the surface layer no longer exists where the 
drops touch. The tension on the remaining surface layer forces 
the drops together into a new sphere, since the latter has the 
smallest surface of any shape which a given volume of substance 
may take. 

The attraction of like particles for each other is called cohe- 
sion, in contrast to adhesion, the attraction between unlike 
particles. 

If a clean glass plate (free from grease) is allowed to touch 
the surface of water and is then gently raised, the water will be 
lifted with it several millimeters before the surface breaks and 
the glass is released. The water is said to adhere to the glass. 
As a matter of fact the water molecules hold tighter 
to the glass than to each other, for when the separa- 
tion comes water parts from water and not from 
glass. The surface of water in a small, clean glass 
vessel is concave upward because of the spreading of 
the water upon the surface of the latter. If a glass 
capillary tube is placed in water the latter rises in 
the tube high above its level outside (Fig. 103.4). Fig. 103.4 
The rising of the water in the tube against the force 
of gravity occurs because of the attraction of the glass for the 
water molecules, and stops when this force is just balanced by 
the weight of water. 

On the same principle, when blotting paper or filter paper is 
put into water the latter rises in the paper even against the force 
of gravity. Before blotting paper was introduced, ink was 
dried by dusting sand over it. The excess ink was taken up 
(adsorbed) on the surface of the sand grains and was removed 
with the latter. The adsorption of liquids by fine powders has 
become of great importance in industry. Nitroglycerine is 
adsorbed by a very finely divided silica (silicon dioxide) called 
infusorial earth, an important substance which will be spoken 




472 



Introduction to General Chemistry 



of again (732). In this form, called dynamite (692), nitro- 
glycerine is much more safely handled than as a pure liquid. 

All liquids do not wet (adhere to) all solids. For example, 
mercury does not wet glass. When mercury is contained in a 
glass vessel the shape of its surface is convex upward 
(Fig. 103.8). This curvature is due to the usual 
I pulling inward of the molecules of the surface layer 

j of the liquid. If a glass capillary tube is thrust 

J into mercury, the level of the latter inside the tube 

i—i ^J is lower than outside simply because the pull inward 
HfiiH of the surface layer resists the deformation made 
Fig. 1035 by the tube. Mercury is not adsorbed by blotting 
paper. Water, of course, behaves the same way 
toward surfaces which it does not wet. Apparently the force 
of adhesion depends on the nature of the substances in question. 
727. Surface Areas. — The foregoing examples are sufficient 
to illustrate well-known surface phenomena. Table XXXVI, 
which gives the surface area of a cube with continuous subdivi- 
sion, illustrates how greatly the surface of a given mass of sub- 
stance may vary, and consequently how effects which are 
ordinarily slight may become of great importance. 



TABLE XXXVI 



Length of Edge of Cube 


Number of 
Cubes 


Total Surface 


1 cm 

1 mm 


I 
IO 3 
IO 12 
IO 21 


6 sq. cm. 
60 sq. cm. 
6 sq. m. 
6000 sq. m. 
» 


. 001 mm 


0.000,001 mm 



If these areas are changed into the system to which we are 
accustomed, we find that if a cube one centimeter on one side 
is divided into little cubes, each o. 001 mm. on one side, the 
total surface is about 65 square feet. If, however, the original 
cube is subdivided into smaller cubes, each 0.000,001 mm. on 
a side, the total surface- is almost an acre and a half! 

728. Adsorption of Gases by Charcoal. — If charcoal is 
freshly heated and thrust into ammonia gas confined over 



Disperse Systems 



473 



mercury, the mercury will be seen to rise in the tube as the char- 
coal cools and finally to remain high above its old level (Fig. 104), 
showing that the ammonia has been taken up by the charcoal. 

There remains in the latter the 
cellular structure of the sub- 
stance from which it was made, 
so that it is composed of innu- 
merable tiny pores and therefore 
presents an enormous surface. 

H I H The fast-moving ammonia mole- 

, '■.;■■ cules encounter this and are 

" held upon it. A few of the 

Fig. 104 iiTi i 

molecules are able to escape the 

carbon, apparently, for all are not taken up, and finally there is 

an equilibrium between the free ammonia and that adhering 

to the charcoal. The ammonia is said to be adsorbed by the 

charcoal. The higher the temperature, the more rapidly is 

equilibrium reached, but the smaller is the amount adsorbed. 

It was on this account that the charcoal was freshly heated 

before it was used in the foregoing experiment, in order that its 

surface should be free from other gases. Charcoal is a very 

good adsorbent for gases, but it does not adsorb all gases to an 

equal degree. Thus a given sample was found to adsorb 90 

times its own volume of ammonia, 35 times its volume of carbon 

dioxide, but only 1 . 7 times its volume of hydrogen. In the 

recent war, filters of specially prepared cocoanut-shell charcoal, 

mixed with other chemicals, were used in gas masks to remove 

poison gases from air inhaled through them. 

729. Air and Glass. — So strongly is air adsorbed on the sur- 
face of glass that great difficulty is found in preparing barometer 
tubes which shall have a perfect vacuum in the space over the 
mercury. It is not sufficient to fill a glass tube with mercury and 
then invert it with the open end under mercury. The tube is 
filled, put under vacuum, and heated for some time before the 
air is removed and the tube is ready to be put in place. 

730. Water Vapor and Glass. — Accurate workers have long 
known that glass surfaces are covered with a thin water him 



474 Introduction to General Chemistry 

which is with difficulty removed. Water molecules encounter 
the glass surface and are held there owing to the very strong 
attraction of water to glass, to which attention has already been 
called. 

These and other experiences show that on the surface of every 
solid exposed to a gas we may expect to find an adsorbed layer 
of the latter. If the solids are finely divided, so that the surface 
is large, the result may become of great importance. 

731. Adsorption and Catalysis. — We have many times noticed 
the catalytic effect of finely divided metals on gas reactions. 
For example, the union of hydrogen and oxygen (303) is greatly 
accelerated by the presence of finely divided platinum. We 
find that platinum takes up both hydrogen and oxygen in con- 
siderable quantity. As early as 1844 Faraday pointed out that 
one result of the presence of a metallic catalyst is that on its 
surface the gases undergoing reaction are at far greater con- 
centration than in the gas mixture itself. Such a layer would 
react far more rapidly than the main mixture. Undoubtedly 
adsorption is the first stage of the reaction on these contact 
agents. But some absorption, or penetration of the adsorbed 
gas beneath the surface on which it is at first held, follows, and 
in some cases compounds are formed. 

732. Adsorption from Solution. — If a little of the dyestuff 
methyl violet, often used as the purple tint of indelible ink 
pencils, is added to water, a beautiful purple solution results 
which may be filtered unchanged. But if charcoal (from sugar, 
for instance) is shaken in this solution and the mixture is then 
poured on the filter, the filtrate is found to be colorless. The 
methyl violet has been adsorbed by the charcoal. That the 
dye has not been destroyed may be shown by pouring alcohol 
through the filter. The filtrate again shows the brilliant color 
of the dye. An enormous number of substances are adsorbed 
from water by charcoal. It is because of this property that it is 
an effective filter for purifying water. However, there are limits 
to the amount that a given column of charcoal can adsorb, so 
that it soon loses this power. Bone black, charcoal made from 
bones, is extensively used to take out objectionable impurities 



Disperse Systems 475 

from sugar solution in the process of refining sugar. Fusel oil, 
a poisonous by-product present in crude whiskey, is also removed 
from the latter by filtration through charcoal. 

Besides charcoal there are other good adsorbents. Fullers 
earth, a very fine clay (mainly aluminum silicate) which varies 
in composition, is an important industrial adsorbent for the 
purification of edible oils. The particles of this earth are usually 
between 0.007 °f a millimeter and 0.00021 of a millimeter in 
diameter. Another important adsorbent is infusorial earth or 
kieselguhr, a deposit of the skeletons of diatoms, which are tiny 
aquatic organisms. Beds of this substance of as much as a 
thousand feet in thickness are found in the United States. Other 
adsorbents are finely divided metals, plant and animal fibers 
such as cotton, silk, wool, etc. As has been pointed out in the 
case of metals (731), absorption may of course follow adsorp- 
tion in the action of these substances. The fundamental require- 
ment for a good adsorbent seems to be an enormous surface of 
contact with the solution. 

733. Suspensions Produced by Grinding of Solids under 
Liquids. — If clay is stirred up in water a turbid mixture results. 
First the coarser particles settle to the bottom, and then gradually 
finer and finer particles follow as time goes on. Pulverized 
emery (174), used for grinding, is graded according to the num- 
ber of minutes required for it to settle after being stirred in 
water. That which will settle in one-half minute is a coarse 
grade. Ten-minute emery is for very delicate work. Careful 
grinding of any substance will produce powders which are still 
slower in settling than any of these, but we have other ways of 
producing them. 

734. Arsenious Sulfide Suspension.— Arsenious sulfide, As 2 S 3 , 
is very insoluble in water. If this solid is shaken up in pure 
water only a minute trace will be found in a liter of the latter 
after filtration. When arsenious acid, H 3 As0 3 , mixed with a 
little hydrochloric acid, is treated with hydrogen sulfide, a pre- 
cipitate of arsenious sulfide appears. But if the hydrochloric acid 
is omitted, only a yellow, opalescent liquid results. When the 
latter is poured through filter paper, merely a trace of precipitate 



476 Introduction to General Chemistry 

is held back and the liquid passes through unchanged. That 
the latter is not a supersaturated solution may be easily estab- 
lished by adding to it a little of the solid arsenious sulfide. No 
settling out of the arsenious sulfide follows, as would be the 
case if the solution were supersaturated (123). 

In whatever form the arsenious sulfide exists in the yellow 
liquid, its presence in water seems to have only a slight effect 
upon the boiling-point and freezing-point of the latter. The 
osmotic pressure (711) of such a solution is very small. If we 
calculate from the latter the amount of arsenious sulfide neces- 
sary to give one gram molecular weight, we obtain enormous 
numbers; for instance, in one case six thousand grams was the 
result, a number which is more than twenty-four times that 
indicated by the formula As 2 S 3 . 

The examination of the liquid under the ordinary microscope 
shows nothing. But if it is examined even with a crude form 
of the ultra-microscope described in 706, it is found to be full of 
dancing particles which must be arsenious sulfide. All the 
evidence shows that the arsenious sulfide exists in the form of 
particles which are very large compared to the simple arsenious 
sulfide molecules, but small compared to particles which settle 
from solution. 

735. Colloids. — Such a non-settling suspension is referred to 
as a colloidal solution, or as a colloidal suspension. Usually 
matter is said to be in a colloidal state if it is too finely divided 
to be held back by a good filter paper and still coarse enough 
to be seen in the ultra-microscope. This means that the average 
diameter of the particles is between 0.000,1 and 0.000,001 mm. 
The setting aside of these systems as a special class is obviously 
entirely a matter of convenience. Colloidal suspensions of 
substances which are solids when in massive form are often called 
suspensoids, while colloidal suspensions of liquids are called 
emulsoids. The name colloid, which means glue-like, was 
originated by Thomas Graham, the first important investigator 
of this subject. He worked mainly with gum arabic, starch, 
glue, and glue-like substances which belong to a more complex 
type than we have yet studied. 



Disperse Systems 



477 



736. Properties of Colloidal Arsenious Sulfide, Diffusion.— 

The most striking thing about the colloidal arsenious sulfide 
suspension is its stability. If it is kept in good-grade glass 
bottles, months may pass without its settling. 

If two test tubes are half filled with 5 per cent hot gelatin 
solution which is allowed to cool and set to a gel, and to one tube 
is added a layer of copper sulfate solution, while to the other the 
yellow arsenious sulfide suspension is added, after about twenty- 
four hours the blue solution will have penetrated well into the 
gel layer, but the yellow solution will not have entered the layer 
below it. The effect can best be seen by corking the tubes and 
then inverting them. The gel in the copper sulfate tube will be 
found to be partly colored, while the other gelatin layer will 
be found to be unchanged. Apparently the rate of diffusion 
of the arsenious sulfide particles is very slow. This is explained 
by the work of Perrin (707). Each arsenious sulfide particle 
is very much larger than a copper sulfate molecule or a copper 
ion, and since it is at the same temperature its velocity must be 
much smaller. 

737. Effect of an Electrical Current on Colloidal Arsenious 
Sulfide. — Colloidal arsenious sulfide solution is a very poor 
conductor of electricity, but if we fill a U-tube 

with this material and pass a no-volt current 
through it, in from ten to twenty minutes a 
colorless layer will be plainly visible at the 
negative electrode, and the region near the 
positive electrode will be deepened in color. 

The effect is better shown by using the 
device described by Professor A. A. Noyes 
and shown in Fig. 105. The ends of the 
inner tube are covered with thin parchment 
paper, or better with goldbeater's skin. Over 
the ends are fitted extension tubes, which are 
joined to the U-tube by rubber. The device is 
inverted for filling through the hole, which is finally closed by 
sliding a rubber tube, H, over it. The spaces around the 
electrodes are filled with ordinary distilled water. The shaded 




Fig. 105 



478 Introduction to General Chemistry 

area of the figure shows the region of yellow color after the 
current has been passing for some time. 

Since the suspension drifts away from the negative and 
toward the positive electrode, its particles are negatively 
charged. We have already noticed that when unlike substances 
are in contact, a loss of electrons by one and a gain of electrons 
by the other are very likely to occur (474). It would not be 
surprising if this were the cause of the existence of this difference 
of potential between the solid and the liquid. But we shall 
soon see that other factors may also account for this phenomenon. 

738. Effect of Electrolytes upon Colloidal Arsenious Sul- 
fide. — We have already seen that in the presence of hydrochloric 
acid the colloid does not form, and we find 'that if a little hydro- 
chloric acid is added to the suspension the liquid becomes very 
much more turbid, and with further addition of the acid a pre- 
cipitate soon appears and settles to the bottom of the container. 
The curious thing is that any other good electrolyte will accom- 
plish the same thing. 

The positive ion seems to be the active part of the electrolyte; 
for about the same concentrations of salts which have univalent 
positive ions are necessary to cause complete coagulation of the 
precipitate from a given volume of solution. Smaller concen- 
trations of salts which have bivalent positive ions are needed, 
and very much smaller concentrations of salts which have 
trivalent positive ions. Relatively little difference is made by 
varying the negative ion of the salt. 

739. Adsorption of the Precipitating Agent. — Careful experi- 
ments have shown that the active ion is carried down by the 
precipitate, leaving in solution an equivalent amount of the 
corresponding ion of water; thus arsenious sulfide coagulated 
with calcium chloride contained calcium, and the liquid left 
behind contained an equivalent concentration of hydrogen 
chloride. 

Apparently the positive ion neutralizes the charge on the 
suspension and is carried down by the precipitate. Since the 
proportion of ion carried down by .the precipitate bears no simple 
and constant relation to the weight of the precipitate, as would 



Disperse Systems 479 

be the case if a chemical compound of the type we know had 
been formed, it is apparent that the product is of a different 
class. For the sake of convenience we may call the former an 
adsorption compound. If the adsorption compound is washed 
with pure water, the calcium is not removed, but it can be taken 
away by washing the substance with some other electrolyte, 
ammonium chloride, for example. In this case ammonium 
takes the place of calcium with the arsenious sulfide, and the 
calcium is found as calcium chloride in solution. 

The proportion of the adsorbed ion to the precipitate is 
usually very small, though of course different with different prep- 
arations according to the valence of the ion, the charge on the 
suspensions, and other conditions. For example, in the case of 
certain colloidal arsenious sulfide suspensions coagulated with 
calcium chloride the proportion was found to be one equivalent 
of calcium to fifty equivalents of the sulfide. If very pure sub- 
stances are desired, as is the case in exact analysis, even this 
small proportion may be of moment, especially if the adsorbed 
ion has a high molecular weight. 

Experience shows that the adsorption of ions from solution 
by a precipitate in the processes of its formation is the rule and 
always must be taken into consideration in the preparation of 
pure substances. 

740. Influence of the Charge on the Stability of a Suspen- 
soid. — If in the experiment on the migration of the suspensoid 
(737) the particles of the latter had been allowed to reach the 
positive electrode, they would have been precipitated. Appar- 
ently the existence of the charge on the suspension is necessary 
for its stability. We can understand that the presence of like 
charges on the arsenious sulfide would have a tendency to keep 
them from coming together, but undoubtedly the effect of the 
charge is more complex than this. 

741. Preparation of Colloidal Ferric Hydroxide. — Another 
fnethod of preparing colloidal solutions is illustrated in the follow- 
ing preparation of colloidal ferric hydroxide. Ferric chloride 
solution, which is of course acid in reaction (436), is treated with 
ammonium carbonate as long as the precipitate which first 



480 



Introduction to General Chemistry 



forms redissolves. The dark-red solution thus made contains 
mainly ferric chloride, colloidal ferric hydroxide, and ammonium 
chloride. The ferric hydroxide and ammonium chloride are 
produced as the ammonium carbonate 
reacts with the hydrochloric acid formed 
by the hydrolysis of ferric chloride : 

FeCl 3 +3H 2 O^Fe(OH) 3 +3HCL 



Fig. 106 



This mixture is next placed in a parchment- 
paper bag (Fig. 106), and the whole is sus- 

pended in a water bath through which a 

current of fresh water is constantly flow- 
ing. If this arrangement is left for about 
four days the contents of the bag will be found to give only 
a very small test for chloride, showing that virtually all of the 
ferric chloride has been transformed into ferric hydroxide. 
The preparation is in fact colloidal ferric hydroxide. 

742. Properties of Colloidal Ferric Hydroxide. — When the 
liquid is poured through filter paper no precipitate is left on the 
latter. The freezing-point, boiling-point, and osmotic-pressure 
determinations all show that the substance is present in the 
form of particles which are large in comparison with simple 
molecules. The dark-red liquid is stable if carefully kept. If 
it is placed in the apparatus, as shown in Fig. 104, and an electric 
current is applied, the red substance is found to migrate toward 
the* negative electrode, proving that this suspension is positive 
to its solution. 

Electrolytes precipitate the red hydroxide, leaving the liquid 
colorless. This time it is the negative ion of the electrolyte 
which is active and goes to form the adsorption compound. 
Again the valence of the precipitating ion is found to be an 
important factor. In a given instance about one-fortieth as 
great a concentration of sulfate as chloride ion was needed to 
coagulate a given amount of ferric hydroxide. 

When the latter suspension is added a little at . a time to 
arsenious sulfide suspensions, the mixture becomes turbid, and 
finally a precipitate of the two substances settles out. This is 



Disperse Systems 481 

the result of mutual adsorption and precipitation. Curiously 
enough the addition of a large excess of either colloid does not 
give a precipitate, but gives a complex colloid. 

743. Explanation of the Preparation of Colloidal Ferric 
Hydroxide.— If a layer of the colloidal ferric hydroxide is put 
over a layer of gel, as was described in the case of arsenious 
sulfide (736), the ferric hydroxide will be found to diffuse very 
slowly indeed. This fact was made use of in the original prepara- 
tion of the colloidal. The ferric hydroxide made by the action 
of ammonium carbonate did not diffuse through the parchment 
paper with appreciable speed, and at the same time it adsorbed 
the ferric chloride so that the latter remained in the bag. The 
hydrochloric acid and the ammonium chloride, however, passed 
readily through the paper and were washed away on the other 
side. The continuous loss of hydrochloric acid allowed the 
hydrolysis of the ferric chloride to go to completion. The 
separation of colloids from dissolved substances by a process 
of diffusion through a membrane is called dialysis and the 
apparatus is called a dialyzer. 

744. Colloidal Silver.— If an electric arc is passed between 
two silver wires submerged in water, a dark cloud will form 
around the electrodes. This is, in fact, a colloidal suspension 
of silver. A little alkali added to the liquid will increase the 
stability of the suspension. The dark liquid has the same general 
properties as those of the other suspensions we have described. 
It is electro-negative to water. This method of producing a 
colloidal suspension of a metal by volatilizing the latter in the 
electric arc is named Bredig's method after its discoverer, the 
German chemist Bredig. 

745. Protecting Agents. — The silver suspension can be made 
of considerable concentration if about 1 per cent of gelatin is 
added to the solution. The effect of the gelatin may be shown 
by mixing silver nitrate (N/10) and hydrochloric acid (N 10) 
solutions, to each of .which about 1 per cent of gelatin has been 
freshly added. Instead of the copious white precipitate which 
we have so often seen result from mixing these reagents, only a 
slight white turbity appears. The gelatin has prevented large 



482 Introduction to General Chemistry 

particles from forming. Many other agents besides gelatin 
stabilize suspensions. Their function is not understood, but it is 
thought that they surround the particles, thus preventing further 
union between them. Hence they are called protecting agents. 

746. Red Gold Suspensions. — A very useful reagent in 
making suspensions of metals is tannin, a complex organic sub- 
stance which is both a good reducing agent and a protecting 
agent. A very finely divided gold suspension may be made 
with a neutralized (about 1 per cent) gold chloride solution and 
dilute (o . 1 per cent) solution of tannin, according to a method 
described by Dr. Wolfgang Ostwald. First a few drops of the 
gold solution are mixed with 100 c.c. of water, then a few drops 
of the tannin solution are added, and the mixture is heated for 
a few minutes, with constant shaking. Meantime the red color 
of the gold suspensoid appears. More gold chloride and tannin 
may be added alternately a little at a time. 

The particles of the red gold suspension are very tiny, usually 
about one or two one-hundred-thousandths of a millimeter in 
diameter. These were the particles described by Zsigmondy 
(70.6). They are negative to the water in which they are 
suspended and are precipitated by positively charged ions. 
Coarser suspensions of gold may be violet or blue in color. 
Brown gold suspensions settle in a very short time. 

747. Summary of Work with Suspensoids. — The general 
methods of preparing suspensoids are illustrated in the foregoing. 
Either the material is subdivided by grinding or volatilization 
in the electric arc, or it is formed in solution and the particles 
are not permitted to grow to a size large enough to settle from 
the liquid. The absence of electrolytes and the presence of pro- 
tecting agents assist in making concentrated suspensions. 
Colloidal suspensions of an enormous number of different sub- 
stances have been prepared. Apparently any substance can 
exist in the colloidal state in liquids in which it is not soluble. 
Thus colloidal sodium suspension has been prepared in ether, 
phosphorus in water, sodium chloride in benzene, etc. 

Apparently many inorganic substances stabilize colloidal 
suspensions to a moderate degree, for we often see opalescent 



Disperse Systems 483 

liquids form when precipitation in dilute solutions is in progress. 
As the last of the precipitating agent is added, the solution be- 
comes clear and the precipitate settles. Thus if 5 c.c. of N/20 
silver nitrate is added to 20 c.c. of N/10 potassium chloride solu- 
tion, a part of the silver chloride appears as a non-settling 
suspension. If more and more silver nitrate is added, with 
constant shaking of the mixture, the liquid finally becomes 
transparent and the precipitate settles just as the amount of 
silver nitrate added becomes nearly equivalent to the potassium 
chloride. The excess potassium chloride present at first stabi- 
lizes the silver chloride suspension. 

The origin of the charges on these suspensions is not under- 
stood. As has been said, it would not be surprising to find a 
frictional charge on them, but we must also consider that an 
unequal adsorption of positive and negative ions from solution 
might also account for the existence of these charges. In addi- 
tion, the loss of ions from the suspended particles to the solution 
might occur. The fact that all suspensions of bases are positive 
to the liquid in which they exist would seem to favor the idea 
of the ionization of the solid particles in these cases. Thus the 
hydroxide particle might lose one or more hydroxyl ions to the 
solution, leaving a residue which would be a very large positive 
ion. But in explaining the presence of these charges we must 
also consider that most other suspensions in water are negative. 
Thus, finally, suspended clay, lamp black, metals, sulfur, salts, 
etc., are negative to their solutions as a rule, though not always. 
Very probably all three possible causes cited for the existence 
of the charges come into play at different times. 

748. Tyndall Effect. — When a pencil of light from a lantern 
is brought to bear on water, we see only a faint glow over its 
path through the latter; but as soon as a colloidal suspension 
is added, the path of the light ray appears as a cone of bright 
cloud. This effect is called the Tyndall effect in honor of 
J. Tyndall, who was the first to make extensive use of this 
phenomenon. The Tyndall effect is, of course, best observed 
in a darkened room. It depends on the difference in refracting 
power for light of the particles of the suspension and of the pure 



484 Introduction to General Chemistry 

liquid. In some cases this difference is not great, so that no 
cone appears. Hence the absence of the Tyndall cone does not 
always mean that the liquid in question contains no colloidal 
matter. The ultra-microscope is of course simply a refined 
apparatus for examining the Tyndall effect in detail. It is 
interesting to note that concentrated solutions of sugar, sodium 
acetate, and many other salts show the Tyndall effect. 

749. Test for the Charge on a Suspension. — A crude distinc- 
tion may be made between positive and negative colloids as 
follows: Strips of filter paper are suspended so that one end of 
each strip is wet by a colloid. If a strip dips into an arsenious 
sulfide or gold solution, the liquid which rises in the paper will 
show the color of the colloid. The colloid is spreading through 
the paper, though usually not so rapidly as the water. But if 
the strip dips into a ferric hydroxide solution, the ferric hydroxide 
will be found to diffuse but a little way into the paper, although 
water from the suspension rises through the latter as easily as 
in the other cases. The paper is negative to water, and so the 
positive colloid is coagulated as it attempts to spread through it, 
while the negative colloid is not interfered with. If turbid 
suspensions are to be tested, they should be filtered before the 
trial is made, otherwise coarse particles will impede the rise of 
the colloid through the paper. 

750. Important Suspensoids. — Suspensoids have become of 
considerable practical importance. Colloidal silver is an 
important antiseptic. It is prepared with a protecting agent 
so that it may be sold as a dry powder (argyrol), which forms a 
dark-brown colloid upon the addition of water. Colloidal copper, 
mercury, and sulfur have also come into use in medicine. India 
ink is mainly colloidal carbon. The important lubricants aquadag 
and oildag are colloidal graphite in water and oil, respectively, 
each protected by tannin. The lubricating power of graphite is 
much improved when it is in the colloidal state. 

We shall next take up emulsoids. These, as we shall see, are 
of tremendous importance in biochemistry. 

751. Oil and Water .—If oil, benzene for example, is shaken 
up in water, for a few seconds there is a general mixture of oil 



Disperse Systems 485 

and water. Then two cloudy layers separate. The lower is an 
emulsion of benzene in water, and the upper an emulsion of 
water in benzene. Very soon, however, the two layers become 
transparent. Momentarily the oil and water are in the colloidal 
state, but this is not permanent. If a drop of water touches 
another drop of water, the two promptly coalesce (726). The 
oil has very little attraction for the water, and its drops continue 
uniting just as the water drops do until two layers are made from 
all the small drops. If, however, soap (678) is added and the oil 
and water layers are shaken together, a stable suspension of oil 
in water will form. The suspension may be termed an emulsoid 
if the droplets are small enough (735). 

752. Cleansing Action of Soap. — Apparently the emulsifying 
power of soap is an important factor in its cleansing or detergent 
action. Thus a little fine dirt shaken with water settles out, 
but shaken with soap and water it remains in suspension. Dirt 
usually sticks to the soap solution more than it does to the 
fabric being washed, and hence rinsing carries off the dirt. 
Undoubtedly some hydrolysis of soap in water does occur, since 
soaps are all salts of the very weak fatty acids (678), but the 
alkali formed is not an important factor in the detergent action 
of soap, since alkali alone has no such power to emulsify mineral 
oils, etc. Vegetable oils can be emulsified to some extent with 
dilute alkalies. But it must be remembered that these are 
esters (677) and that some soap is formed by the interaction of 
esters with the alkali. If alkali, washing soda, or ammonia is 
added to soap solution, its detergent power is increased. Of 
course the hydrolysis of soap is decreased at the same time. 

753. Nature of Soap Solution. — Ultra-microscopic examina- 
tion of soap solution is unsatisfactory, presumably because there 
is very little difference in the power of soap and water to refract 
light. The fact that soap has little effect on the boiling-point 
of water seems to indicate a very high molecular weight. All 
evidence points to the conclusion that the soap molecules in 
solution are groups of large numbers of simple molecules. 

754. Work of Harkins and of Langmuir. — It has long been 
known that a soap solution is more concentrated in the surface 



486 Introduction to General Chemistry 

layer than in the inner region. Professor W. D. Harkins and 
co-workers have found that this is true of solutions of sodium 
oleate (678, 681) as dilute as 0.002 normal. The surface layer 
of a 0.002 normal solution is saturated. Stronger solutions 
up to 0.1 normal have the same concentration in the surface 
layer as have these very dilute solutions. In each case there is 
equilibrium between molecules of soap in the surface layer and 
in the body of the solution. Evidently the surface layer allows 
the escape of relatively few molecules to the solution, or the 
inequality of concentration could not be kept up. Further, 
Professor Harkins and Dr. Langmuir, of the General Electric 
Company, working independently and by different methods, 
have shown that the soap molecules of the surface layer are not 
simply jumbled together but are arranged in a definite order, 
The long sodium oleate molecules 

OHHHHHHHHHHHHHHHHH 

I! I I I I I I I I 1 I I I I I I ! ! 

Na-0-C-C.C-C-C-C-C-C.C = C-C-C.C.C-C-C.C.C.H 

I I I I ! I I I I I I I I I I 

HHHHHHH HHHHHHHH 

O 

II 
are all placed with the (Na • O • C • ) group toward the body of the 

solution and the long hydrocarbon chain (664) outward. 

It is an old and well-known rule that "like dissolves like." 

Thus oils are good solvents for fats and greases but poor solvents 

for sugar, while water is a good solvent for the latter but a poor 

O 

solvent for the former. Apparently water attracts the Na • • C 
group of the soap molecule strongly but has no attraction for the 
long hydrocarbon chain. As a result of this attraction on one 
end of the long molecule, this end is pulled in toward the liquid, 
and the hydrocarbon chain is left on the surface. The surface 
of the soap solution is therefore that of an oil, and it is on this 
account that soap solutions can wet oils while water cannot. 

755. Other Emulsifying Agents. — Other emulsifying agents 
are in common use, but their function is not so well understood. 



Disperse Systems 487 

Thus in mayonnaise dressing the yolk of an egg is used to keep 
the oil and vinegar together. Many substances are capable of 
forming emulsoids without the assistance of special agents. 
Examples of these are gelatin, gum arabic, egg white, etc. These 
are all substances which are themselves wet by water. 

756. General Properties of Emulsoids. — Emulsoids are 
sharply differentiated from suspensoids by their relatively high 
viscosity. The tiny particles of both show Brownian move- 
ments (705). Unlike suspensoids, many emulsoids seem to be 
without electrical charges. Salts in small amounts have little 
effect on these systems. In large amounts they may cause 
precipitation. Thus soaps are precipitated (" salted out") from 
solution upon the addition of considerable quantities of sodium 
chloride. Although we have discussed emulsions in water only, 
it is obvious that they may exist in other media. 

757. Important Emulsoids. — Virtually all fluids of plant and 
animal bodies are emulsoids. Thus the sap of plants and the 
blood and milk of animals are complex emulsoids. 

758. Gels. — If the solvent is evaporated from a suspensoid, 
the latter is left as a powder. But if the solvent is evaporated 
from an emulsoid, or if a concentrated emulsoid is cooled, a gel 
usually forms, though not always. Soap melted in hot water 
sets as a gel on cooling. Solutions of gelatin of greater concen- 
tration than 0.25 per cent will set at temperatures above o°. 
These gels are stiffer the higher the concentration of the original 
emulsion. 

Unfortunately they refract light little differently from water, 
so that an examination of their structure under the ultra-micro- 
scope is not satisfactory. Gels have been precipitated by the 
addition of alcohol, etc., and then examined. Their structure 
then seems to be that of a delicate latticework or honeycomb, 
but it is more likely that this form is acquired in the process of 
precipitation. 

We have already seen that substances in solution diffuse 
through gels very readily. In most cases they do so as easily 
as through water, but colloidal substances, either emulsoid or 
suspensoid, diffuse through gels only very slowly. Some gels 



Introduction to General Chemistry 



are elastic and others are not. If the former are stretched they 
become warm, and cool again on contraction. Another very 
striking property of gels is their ability to absorb (731) water. 
If a warm, concentrated gelatin solution is poured out on a glass 
plate and allowed to set, it may be cut into pieces of equal volume. 
The latter should then be dried. If a piece is allowed to soak 
in water, it will be found to swell. If at the same time other 
pieces of the same weight are allowed to soak about twenty-four 
hours in dilute acid or dilute alkali of equivalent concentration, 
the rate of swelling of the gel will be found to be most in the acid 
solution, next in the alkali, and least in the pure water. Salts 
added to pure water may increase or decrease the rate of swelling. 
Still another important property of gels is their separation into 
two layers, one liquid and the other gel. If evaporation is 
prevented, all gels act in this way in the course of time. 

759. Plant and Animal Tissue. — Virtually all of the firmer 
parts of animals and plants are gels of great complexity, but the 
properties given above for gelatin belong to most of these sub- 
stances. For example, the experiment with the gelatin squares 
may be repeated with animal tissue (frogs' legs, for example) 
with the same type of result. The development of a watery 
fluid by a gel is repeated in the formation of many animal and 
plant secretions. 

760. Rubber. — If rubber, which is a typical colloid, is allowed 
to soak in benzene or carbon disulfide, it swells enormously, at 

the same time taking up the liquid. 
Upon being stretched rubber becomes 
warm, and cools when contracting, just 
as do other gels. An interesting appli- 
cation of the general law governing all 
changes, applications of which we have 
observed so many times (367), may be 
shown in the following experiment: A 
stout rubber band is attached by one 
end to a support and is stretched by a 
weight, as in Fig. 107. If a lighted match is held near the 
rubber, so that it is heated quickly but not melted, the weight 



-4 



Fig. 107 



Disperse Systems 489 

will be seen to rise. The application of heat has favored the 
change which absorbs heat, namely, contraction. 

761. Inorganic Gels. — Many inorganic substances form 
gels easily; for example, silicic acid, ferric hydroxide, and 
aluminium hydroxide. These are all extremely insoluble sub- 
stances. Von Weimarn has shown that in general, if a very 
high degree of supersaturation is attained preliminary to pre- 
cipitation, the precipitate will be a gel. Thus barium sulfate 
is usually seen to precipitate as a powder. But if saturated 
solutions of the very soluble salts sodium sulfate and barium 
sulfocyanate are mixed, a gel of barium sulfate forms. Con- 
centrated solutions of sodium carbonate and calcium chloride 
give a gel of calcium carbonate instead of the usual powder. 

762. Relative Stability of Precipitates. — Apparently the 
crystalline form is more stable than either the gel or the powder 
precipitate. On long standing in contact with the solution most 
non-crystalline precipitates become crystalline. The rate of 
change is accelerated by a rise of temperature which increases 
the rate of molecular agitation and usually also increases the 
solubility of the precipitate. Even such substances as gelatin 
can be prepared in crystalline form. A substance can separate 
from solution in the crystalline form only if there are time and 
opportunity for the orderly arrangement of molecules in crystals. 
Hence only the more soluble substances usually appear in this 
form. 

763. Explanation of Adsorption from Solution. — The fact 
that soap is more concentrated in the surface layer than in the 
solution is a phenomenon of a type which is very common among 
substances of complex molecular structure. Most emulsoids 
and suspensoids show this to a marked degree, though substances 
in solution do also to some extent. Under ordinary conditions 
this is not noticeable. But when the surface of the solution in 
question is enormously increased by being mixed with a porous 
or finely divided solid like charcoal or infusorial earth, the net 
result is that a great amount of the colloid goes to the surface 
layer and adheres to the solid with the latter, while the rest of 
the liquid is drained away. The very large extent of the surface 



49° . Introduction to General Chemistry 

means a very great loss of material from the original liquid. 
The latter may be virtually freed from the colloidal matter if 
passed through a thick layer of filter. 

764. Importance of Colloid Chemistry. — The many complex 
systems discussed in this chapter have one thing in common — 
an enormous surface of contact between different simpler 
systems. The porous charcoal and gas, arsenious sulfide and 
water, oil, soap and water, etc., may all be called disperse 
systems. They are also called colloids and their study colloid 
chemistry. Even this brief treatment should be sufficient to 
show the reader the enormous importance of the subject. As a 
matter of fact the majority of practical applications of chem- 
istry involve disperse systems. Students will find the five 
lectures written on this subject for the general pubic by Wolfgang 
Ostwald 1 well worth reading. 

x Ostwald and Fisher, Theoretical and Applied Chemistry. Published by 
John Wiley & Sons, New York. 



CHAPTER XXIX 
THE ATMOSPHERE AND RELATED TOPICS 

765. The Composition of the Air. — We have already learned 
that the most important components of air are nitrogen, oxygen, 
water vapor, and carbon dioxide. In addition to these four, 
there is but one other component the proportion of which exceeds 
0.01 per cent. This is the element argon (513), discovered in 
the air in 1894. The percentage by volume of the five named 
components is given in Table XXXVII. 

TABLE XXXVII 

Percentage Composition or the Air 
by Volume 

Nitrogen 77. 10 

Oxygen . . ..... 20.70 

Argon . . . .'•'.. . . . 0.80 

Carbon dioxide o . 03 

Water vapor (about) . . . 1.35 

Sum 99 .98 

It is self-evident that the air must also contain minute amounts 
of several other gases and vapors, since these are being poured 
into the air from numerous industrial sources and are also being 
formed by natural chemical changes. Among the minor com- 
ponents of the air there may be mentioned hydrogen, methane, 
sulfur dioxide, hydrogen sulfide, ammonia, and nitrogen tetroxide. 
These chemically active gases are continuously being removed 
from the air by various means, so that they never accumulate 
in appreciable proportions. Other minor gaseous components 
of the air will be considered later (791-799). 

766. Why the Composition of the Air Remains Constant. — 
It is plain that oxygen is removed from the air on every hand 
by the burning of substances and by the respiration of all animals. 
If oxygen were not being constantly renewed the percentage of 

4QI 



492 Introduction to General Chemistry 

this element would in time steadily decrease. But as we have 
learned (691), all growing plants take in carbon dioxide and 
water, from which they form, in addition to other products, 
cellulose, or starch, and oxygen, thus: 

6CO 2 +sH 2 O->C 6 H I0 O 5 +6O 2 

The oxygen supplied to the air in this way serves to keep the 
percentage of this element fairly constant. 

Closely connected with the oxygen balance in the air is that 
of carbon dioxide. The removal of this gas by growing plants 
is compensated by its formation in the oxidation of wood, coal, 
fuel gas, and other carbon compounds (356-365). One other 
agency has been of much importance in past geological ages in 
diminishing the carbon dioxide content of the air. This is the 
formation of the shells of aquatic animals, consisting largely of 
calcium carbonate. Immense deposits of limestone (150) have 
been built up from the shells of marine animals. 

The water-vapor content of the air varies greatly from place 
to place and from time to time in a given place. Over the ocean 
the lower layers of the air are nearly saturated with water vapor, 
so that the vapor pressure tends to approach, at each prevailing 
temperature, the value shown in Table VII (112). In desert 
regions the relative degree of saturation (the humidity) is very 
small. 

The percentages of nitrogen, oxygen, and argon in air freed 
from water vapor and carbon dioxide are practically constant at 
all times the world over. The carbon dioxide content rarely falls 
below 0.03 per cent but may reach or exceed 0.04 per cent in 
congested parts of cities. Indoors it may at times run much 
higher. 

767. Dew and Frost. — After sundown the temperature of 
the air falls, on clear nights in particular, because of the more 
rapid radiation of heat in the absence of clouds. If the humidity 
has been rather high by day the air may become supersaturated 
with water vapor at night. For every temperature the vapor 
of liquid water exerts a definite pressure, its saturation vapor 
pressure (112, Table VII). If air containing a fixed proportion 



The Atmosphere and Related Topics 



493 



of vapor falls in temperature below the point at which it becomes 
saturated, the excess water separates out on all exposed objects 
as dew. The dew point is defined as the temperature to which 
a given sample of moist air must be cooled just to reach a condi- 
tion of saturation. If cooled further it forms dew. 

When the dew point lies below zero, if the temperature falls 
sufficiently low, the water vapor deposits in the form of frost 
(ice). 

768. Dust in the Air and Cloud Formation. — The presence 
of dust in the air is, in general, considered nothing less than a 
nuisance; but, as we shall point out shortly, it has a very 
important function in connection with the formation of clouds 
and consequently of rain. For this reason especially we shall 
consider the dust content of the air. 



«=* 




Fig. 108 



A cloud, or a fog which is a cloud at the earth's surface, is 
made up of countless drops of water, each so minute that it 
does not fall with appreciable speed (compare Millikan's experi- 
ment, 467). A cloud may form when a mass of moist air is 
cooled below its dew point. But this is true only under certain 
conditions, as may be demonstrated by a lecture-table experi- 
ment. If air not specially freed from dust, and saturated with 
water vapor, is suddenly cooled several degrees, a cloud or fog 
forms at once; but if the air is entirely free from dust no cloud 
is formed. To show these phenomena experimentally we may 
fill a large bottle (Fig. 108), with air that has been bubbled 



494 Introduction to General Chemistry 

slowly through water in order to insure its saturation with 
vapor. The bottle should be closed with a stopper through which 
passes a glass tube attached by a rubber tube to a second bottle 
half the size of the first. The second bottle should have a two- 
hole stopper and a second tube leading to a vacuum pump. The 
rubber tube connecting the two bottles is now closed by means 
of a clamp and the smaller bottle evacuated. If now the clamp 
is released, the air in the large bottle expands rapidly, partly 
passing over into the smaller bottle, and at the same time a 
dense cloud develops in the larger bottle. The expanding air 
does work in the process, and the energy to do the work is taken 
from the air as heat; in consequence the temperature falls so 
much that the air becomes supersaturated at the lower tempera- 
ture, and the excess moisture separates as a cloud. If we repeat 
this experiment with the single change that the air used in filling 
the large bottle is freed from dust before it is bubbled through 
water, by causing it to filter through a cotton plug {A, Fig. 108), 
no cloud forms upon expansion. If now a small amount of dust 
(for example a little cigarette smoke) is admitted to the partially 
evacuated large bottle a cloud is formed. (The smoke must be 
introduced quickly, before the temperature has time to rise.) 

769. Counting Dust Particles. — Anyone who has noticed the 
particles of dust floating in air, which may be seen when a 
bright beam of light enters a darkened room, would be inclined 
to say that counting the stars or the grains of sand by the seaside 
would be a simple task in comparison with the enumeration 
of the dust particles in a given volume of air. But thanks to 
the facts related in the preceding section the undertaking proved 
rather easy. Without dust, cloud formation does not take 
place. In the presence of dust each dust particle acts as a 
nucleus for the formation of a water drop. Therefore there are 
as many water drops in the cloud as there were dust particles 
present! A practical instrument has been devised by Aitken 
to count, by the aid of a microscope, the drops formed in a known 
volume of air and thus to ascertain the dust content of the air. 
The number of dust particles per c.c. of air varies greatly with 
circumstances, as would be expected. The smallest values, a 



The Atmosphere and Related Topics 495 

few hundred particles per ex., are found in air above mid-ocean. 
Clear mountain air contains a few thousand particles per ex., 
while the air of large cities often contains over 100,000 per c.c. 
In his interesting article in the Encyclopaedia Britannica on 
"Dust," Aitken writes: 

Without atmospheric dust not only would we not have the glorious 
cloud scenery we at present enjoy, but we should have no haze in the 
atmosphere, none of the atmospheric effects that delight the artist. The 
white haze, the blue haze, the tender sunset glows of red, orange and yellow, 
would all be absent, and the moment the sun dipped below the horizon 
the earth would be in darkness; no twilight, no after glows; none of the 
poetry of eventide. Why it may be asked is this so ? Simply because all 
these are due to matter suspended in the air, to dust. 

770. The Ionization of Gases. — The topic about to be dis- 
cussed is related to cloud formation in an important way, and 
for this reason we shall have to digress somewhat before going 
on with our subject. It is well known that air under ordinary 
conditions is an almost perfect electrical insulator. However, 
it may be made appreciably conducting in 
several ways. The phenomena may be 
readily shown by the use of a gold- or 
aluminum-leaf electroscope, Fig. 109. If 
the electroscope has good insulation (sulfur 
or amber, A, Fig. 109) it will retain its 
charge for a long time. If a name of any 

kind is brought near the electrode, B, the -. 

to ^ ' ' Fig. 109 

gold leaf, C, drops almost as rapidly as if 
one had touched B with the finger. The gases in and about 
a flame conduct electricity thousands of times better than 
ordinary air. Elaborate investigations, which cannot profitably 
be discussed here, have proved that part of the gas molecules 
of a flame are electrically charged, half positively, half negatively. 
Consequently the gas is said to be ionized; and just as in the 
case of an ionized solution the gas conducts electricity. 

Every gas can be ionized by a variety of means, among which 
are intense heat, X-rays (476), and radium rays (480). Even 
the most energetic means do not convert into ions more than a 
very small fraction of the total number of gas molecules. 




496 Introduction to General Chemistry 

771. Gaseous Ions and Cloud Formation. — Gaseous ions 
behave like dust particles in being able to serve as nuclei for 
water drops in cloud formation. If dust-free air, saturated with 
water vapor, is ionized by X-rays, by radium, or in any other 
way and is then suddenly cooled (as by expansion in the manner 
described in 768) a cloud will form, each ion acting as a nucleus 
for a single drop of water. 

772. The Formation of Rain. — There are always present in 
the air a small number per c.c. of gaseous ions formed in part at 
least by radioactive matter in the air. These ions, together 
with a much larger number of dust particles, serve as the nuclei 
of the drops forming ordinary clouds. Cloud formation occurs 
when a current of warm, moist air meets a current of cold air. 
If the water drops are large enough they tend to coalesce and 
thus grow so large that they fall to earth as rain. Rain is fre- 
quently accompanied by lightning, the cause of which we may 
now consider. 

773. The Cause of Lightning. — If air is only slightly super- 
saturated with water vapor the negative ions present are much 
more effective in the condensation of moisture than are the 
positive ions, so that it frequently happens that only the nega- 
tive ions are at first removed from the upper layers of air and 
carried to the earth with the falling rain. This results in the 
accumulation of opposite electrical charges in the air and on 
the earth beneath. The lightning that frequently accompanies 
rain is the electric discharge between air and earth, tending to 
neutralize the unlike charges. 

774. The Atmosphere, a Disperse System. — The behavior 
of tiny water and dust particles in the air suggests the behavior 
of small particles in liquids. We have already noted (703) the 
Brownian movement of smoke particles. The accumulation of 
an electrical charge by clouds in the process of formation and 
the formation of rain with the discharge of the former have their 
counterparts in the behavior of suspensoids. The atmosphere 
with its dust and water particles is, of course, a disperse system. 

775. The Liquefaction of Gases. — Faraday's experiments on 
the liquefaction of chlorine have already been described (242). 



The Atmosphere and Related Topics 



497 



We have also seen that ammonia can be liquefied when strongly 
compressed (517). Carbon dioxide gas can be liquefied (633) 
at a very high pressure if at the same time the temperature is 
below 31 . Above this temperature no pressure, however great, 
causes liquefaction. At 31 the pressure required is 72 atmos- 
pheres. At lower temperatures less pressure is required. 
Experiment has shown that for every gas there is a definite 
temperature above which no pressure (however great) will cause 
liquefaction. This temperature is called the critical tempera- 
ture of the gas. The vapor pressure exerted at its critical 
temperature by a liquefied gas is called its critical pressure. 
Any gas may be liquefied if cooled below its critical temperature 

TABLE XXXVIII 



Gas 


Critical Temperature 
in Degrees C. 


Critical Pressure 
in Atmospheres 


Boiling-Point in 
Degrees C. 


Ammonia 

Carbon monoxide 


131° 

— 140 

3T 

52 

- 94 
37 

■ ■ 155 
146 

-243 
-145 
-119 


113 

36 

72 

84 

71 

72 

79 • 

84 

11 

34 

5i 


- 39° 

- 190 
• - 76 

- 35 
-154 

- 88 


Carbon dioxide 


Hydrogen chloride 

Nitric oxide 


Nitrous oxide 


Sulfur dioxide 


— 10 


Chlorine 

Hydrogen 

Nitrogen 

Oxygen 


- 34 

— 252 
-194 
-182 



and subjected to a pressure which need not exceed its critical pressure. 
All known gases except five (hydrogen, nitrogen, oxygen, nitric 
oxide, and carbon monoxide) were liquefied by Faraday by the 
year 1845. The critical temperatures, critical pressures, and 
boiling-points at one atmosphere pressure of a number of gases 
are given in Table XXXVIII. 

Solid carbon dioxide (633) mixed with ether (642) to make 
it a better heat conductor was used by Faraday as a cooling 
agent. By means of this mixture a temperature of — So° is 
easily obtained. Even at — 8o° and at very high pressures 
Faraday could not liquefy the five named gases. These he 
called permanent gases. We now know that these gases differ 
from others only in having critical temperatures considerably 



4Q8 



Introduction to General Chemistry 



lower than — 8o°. They have all been liquefied in more recent 
times. 

776. Liquefaction of Air. — Practical methods for the liquefac- 
tion of air have been devised by Hampson, by Linde, and by 
Claude. The principle of the apparatus of the first two of these 
inventors is the same and will be understood by reference to 
the diagrammatic Fig. no and the following description. Air 
is compressed to 200 atmospheres pressure by a powerful pump, 
A . The work of compression is largely changed to heat so that 
the air becomes very hot. It is next cooled by means of running 
water in a cooler, B, and afterwards passes through a thick- 
walled steel cylinder, C, containing solid caustic potash to free 
it from water vapor and carbon dioxide. The compressed air 





















C 




^v 


B 






/V^\ A 












(rr\- 






Fig. iic 


\\ — 


i 








*w 



next passes through a metal pipe ending in a needle valve, D, 
from which it escapes and expands to normal atmospheric 
pressure. Upon expansion at the needle valve the air falls 
greatly in temperature (compare the preparation of carbon 
dioxide snow, 633), but not sufficiently to cool it to its tempera- 
ture of liquefaction. But the very cold expanded air now flows 
upward through the larger pipe surrounding the pipe leading 
to the needle valve and thus cools the compressed air greatly 
before it escapes. The result is that the air reaching the valve 
becomes colder and colder until finally it reaches its liquefaction 
point; that is, its boiling- temperature at atmospheric pressure. 
From this time on liquid air collects in the lower part of E and 
can be drawn off through a second valve. 

777. The Properties of Liquid Air. — Liquid air is a faintly 
blue mobile liquid. It can be kept in the liquid state at atmos- 




The Atmosphere and Related Topics 499 

pheric pressure only so long as its temperature remains at or 
below its boiling-point, about — 190 . If it is contained in an 
open vessel, it boils away more or less rapidly, and the absorp- 
tion of heat during its evaporation (115) keeps its temperature 
down to its boiling-point at atmospheric pressure. The more 
slowly external heat flows into the liquid air the more slowly it 
boils away and therefore the longer it may be kept. Sir James 
Dewar has invented ingenious containers for liquid air that 
provide the best attainable heat insulation. These are glass 
vessels with double walls, having the space between 
the walls evacuated (Fig. in). Often the walls are 
silvered. Glass is a poor conductor of heat, and as 
there is no air between the walls to transfer heat by 
convection, heat can reach the interior only by con- 
duction, by way of a long path through glass. 
Radiant heat is kept out by the silvered surface, 
which reflects both heat and light. Heat reaches F 
the interior so slowly that a five-liter flask of liquid 
air will lose by evaporation only 20 to 25 per cent of its contents 
in twenty-four hours. In recent years Dewar vessels have come 
into extensive general use (as "Thermos" bottles, " Icy-Hot" 
bottles, etc.). 

Liquid air, being a mixture of liquid nitrogen and liquid 
oxygen, does not have a constant boiling-point. Liquid nitrogen 
boils at —1 94 and liquid oxygen at — 182 . The former, being 
the more volatile, boils away more rapidly, so that the partially 
evaporated liquid is largely oxygen. This is decidedly blue in 
color. Liquid nitrogen is colorless. By a process of systematic 
fractional distillation the component gases of liquid air can be 
nearly completely separated. These separate gases are made 
commercially on a large scale in this way (309, 513). 

778. Liquefaction of Other So-called Permanent Gases. — 
With the liquefaction of air it became apparent that there were 
no permanent gases in the sense that these could not be liquefied. 
It was plain that any gas could be liquefied either at atmospheric 
or higher pressure, provided it could be cooled below its critical 
temperature. Faraday had failed to liquefy rive gases because 



500 Introduction to General Chemistry 

he had no means of cooling them to sufficiently low temperatures. 
With liquid air, boiling at atmospheric pressure, a temperature 
of — 190 is available. Since carbon monoxide boils at — 190 
and nitric oxide at — 1 54 both these gases are readily liquefied 
by cooling them with liquid air. In the case of the former a 
little pressure above atmospheric is required. Hydrogen, how- 
ever, cannot be liquefied at — 190 even when strongly com- 
pressed. This fact indicates that the critical temperature of 
this gas is lower than — 190 . 

779. Liquid Hydrogen. — Hydrogen was condensed to a 
liquid (in appreciable amounts) by Dewar in 1898. The method 
employed was similar .to that used to make liquid air (776); 
but the compressed hydrogen before entering the liquefier, D, 
Fig. no (776), was cooled to — 185 by liquid air. Liquid 
hydrogen is a colorless liquid, about one-seventh as dense as 
water. Its critical temperature is — 243 , its critical pressure 
n atmospheres. It boils at — 252 at atmospheric pressure, 
equal to 21 absolute. 

780. Flames. — That a flame is a burning gas is of course well 

known to the reader; but it may not have occurred to him that 

the flame of a candle or kerosene lamp is a gas flame. In the 

case of a burning candle the wax melted by the heat forms a 

small pool of liquid; this the wick takes up by capillary action 

and brings to the center of the flame, where the 

intense heat decomposes the wax into volatile 

(gaseous) products. By holding a narrow glass 

tube 3 inches long in a candle flame so that the 

lower end of the tube is at the tip of the wick, 

unburned gas may be drawn from the center of 

the flame and burned at the upper end of the 

Fig. i 1 2 tube, Fig. 112. This experiment also shows that 

in the center of the candle flame the gas is as 

yet unburned. It unites with oxygen in the outer layer of the 

flame; this is therefore the hottest part of the flame, while -the 

interior is much cooler. If a piece of writing paper is held for 

a few seconds in a candle flame at the tip of the wick and 

perpendicularly to the latter, it will be scorched in a ring, the 

center of which is unburned. 




The Atmosphere. and Related Topics 501 

The flame of burning wood or coal is formed similarly to 
the candle flame by reason of the preliminary conversion of the 
fuel into gaseous products. 

781. The Bunsen Burner. — If we close the air vents at the 
base of a Bunsen burner the ignited gas burns with a flame 
resembling the candle flame in structure. The flame is luminous 
and is likely to be smoky. If we open the air vents sufficient 
oxygen becomes mixed with the gas to cause much more rapid 
burning, since now it is not necessary for gas and air to mix by 
the rather slow process of diffusion after the gas has left the 
burner tube. With the ideal adjustment of the air vents some- 
what less oxygen is supplied than the total necessary, so that 
some oxygen is taken from the air around the flame. If too 
much air mixes with the gas at the vents, the mixture burns 
with such great rapidity that the speed of ignition exceeds the 
speed with which the mixed gas and air travel upward in the 
burner tube, with the result that the flame 

" strikes back" and "burns at the base." 
In an improved form of burner known as 
the Meker burner (Fig. 113) the design of 
the tube and vents is such that a larger 
proportion of air is taken in than in the 
ordinary Bunsen burner, and air and gas 
are more intimately mixed before reaching 

the top of the burner. To prevent strik- « 

ing back, the wide upper end of the burner 

tube is provided with a metal grid, in the F 

narrow passages of which the hot gases are 

so greatly cooled that the downward speed of the ignition wave 

no longer exceeds that of the upward-moving gas stream. 

The Bunsen flame with open vents consists of two distinct 
parts: the inner cone of a greenish color and the outer cone, 
bluish in color and less luminous. The hottest part of the flame 
is found in the center of the outer cone, just above the apex of 
the inner cone. 

782. Luminous and Non-luminous Flames. — The cause of 
luminosity of a gas flame, such as that from a kk hsh tail" burner, 
has been the subject of extensive investigation. It is probable 



°o° 




502 Introduction to General Chemistry 

that certain compounds in the gas, notably acetylene, C 2 H 2 
(663), are decomposed by heat giving free carbon; and that the 
particles of the latter, being intensely heated, give out light. 
Indeed we have only to hold a cold object in a luminous flame to 
collect on it a deposit of soot (carbon). Furthermore it is well 
known that if acetylene is strongly heated in the absence of air 
it dissociates into hydrogen and carbon. The particles of carbon 
in a luminous flame are finally more or less completely burned. 

The action of air in rendering a Bunsen flame non-luminous 
is said by some chemists to be due to more perfect combustion. 
But this explanation is not quite sufficient, since a gas flame is 
rendered non-luminous by admixture with nitrogen or carbon 
dioxide instead of with air. Probably these inert gases so act 
because the temperature of the flame is reduced by the dilution 
to a point below that at which the acetylene, etc., is decomposed 
before it burns. Air, being four-fifths nitrogen, must also act 
as a diluent in the Bunsen flame. 

We always use the non-luminous Bunsen flame when we 
wish to heat anything most efficiently; and we might be led to 
conclude that a given volume of gas produces more heat as a 
non-luminous than as a luminous flame. This, however, is not 
true ; the heat production is exactly the same in the two cases, that 
is, if the combustion is complete (363). Nevertheless a beaker 
of water will be heated more quickly and a piece of glass will 
be heated hotter in the non-luminous flame. One reason is this: 
the luminous flame radiates heat, as well as light, in much greater 
amount than does the non-luminous flame, as is easily proved 
by holding the hand at a distance of a few inches on one side of a 
Bunsen flame and opening and closing the air vent. Most of 
the radiated heat is lost for practical heating purposes. A* 
second reason why the non-luminous flame is more effective is 
that it is more concentrated (compact), and that the gases are 
in more rapid motion. For both these reasons hot molecules 
strike the object to be heated more frequently than is the case 
in the luminous flame and so raise its temperature more rapidly. 
783. Reactions in the Flame. Bead Tests. — It is perhaps 
a new point of view to consider the flame as a reagent, but that 



The Atmosphere and Related Topics 503 

it is a valuable one may be illustrated by some well-known bead 
tests. If a metaphosphate bead is made in the usual way and 
a tiny speck of a copper salt is melted into it by heating the two 
in the non-luminous flame of the burner, after the resulting bead 
is cool it will be found to be blue, the usual color of dilute solu- 
tions of copper salts. But if the air holes of the burner are partly 
closed, so that a small, luminous sheath appears on the tip of 
the inner cone of the flame and the bead is first melted in this 
sheath and then lowered into the inner cone of unburned gas 
until it is cool, upon removal from the flame the bead will be 
seen to be opaque and reddish, owing to the presence of finely 
divided copper. If the red bead is reheated in the outer flame 
the blue color reappears. As a consequence of such reactions 
the inner cone of the flame is known as the reducing region 
and the outer as the oxidizing region. The reducing region of 
the non-luminous flame is relatively thin, but its depth is in- 
creased by shutting off a little of the air as indicated above. 
Many other bead tests can be made in a similar way. 

784. Colored Flames. — A number of elements if brought into 
a non-luminous Bunsen flame in the form of volatile compounds 

TABLE XXXIX 

Element Color of Flame 

Sodium Yellow 

Potassium ...... Violet 

Lithium Crimson 

Calcium . ' Orange 

Strontium Red 

Barium Green 

Copper Bluish green 

Boron Green 

(usually salts) give to the flame characteristic colors. Thus, 
common salt and other compounds of sodium give intensely 
yellow flames. The flame color is, in general, dependent on the 
elements present irrespective of whether they are free or com- 
bined. Table XXXIX gives a list of the commoner elements 
forming colored flames. 



5°4 



Introduction to General Chemistry 




Fig. 114 



Very interesting and important facts about colored flames 
are brought out by the use of the spectroscope, to the descrip- 
tion of which we shall now turn. 

785. The Spectroscope. — When a narrow beam, A, of white 
light, with parallel rays, strikes a prism, B, in the manner illus- 
trated in Fig. 114, its 
colors are refracted to 
different degrees so that 
the emerging light is 
spread out as a spec- 
trum at C, with its red 

rays least and its violet 

rays most changed in 
direction. A spectro- 
scope (Fig. 115) consists of a prism, A, with one tube, B, 
carrying lenses so arranged as to throw a narrow beam of light 
on the prism, and a second tube, C, carrying a set of lenses 
forming a microscope through which the spectrum is viewed. In 
most cases there is in addition a third tube, D, carrying a scale 
which can be illuminated, by means of which the different colors 
of the spectrum can be located. If one examines with the 
spectroscope the light from an incandescent electric bulb he sees 
the entire spectrum from 
red on the one hand to 
violet on the other, with 
no single color noticeably 
brighter than any other. 
The same kind of spectrum 
is given by any white-hot 
body, as for example an 
incandescent gas mantle. 

Such a spectrum is called a continuous spectrum. 
gas flame also gives a continuous spectrum. 

786. Bright-Line Spectra. — An entirely different picture is 
presented when one looks at a non-luminous Bunsen flame made 
yellow by a sodium salt. Instead of the entire spectrum only a 
bright-yellow line appears. If the instrument is a good one 




Fig. 115 



A luminous 



The Atmosphere and Related Topics 505 

and a very narrow beam of light is used, the yellow line is found 
to consist of two parallel lines very close together in the position 
in the scale occupied by the yellow part of the spectrum when 
the entire spectrum is present. The spectrum of the element 
lithium is still more striking, in that it appears as a single bright 
line of the purest red. The spectrum of potassium shows two 
lines, one in the red but not in the same position as that of lithium 
or of the same shade of red, and another in the violet. Some 
elements like calcium give a rather complex spectrum made of 
several lines, some of which are rather broad. Each element 
gives its own characteristic spectrum, so that if one is familiar 
with the various line spectra of the elements it becomes a very 
simple matter to identify at once any element which gives a 
colored flame. This identification is made easy because each 
line of the spectrum always appears at a fixed position on the 
scale of the instrument, so that the observer has only to note the 
scale position of the lines without considering critically their 
colors. It is not difficult to detect spectroscopically the presence 
of two or more elements in a mixture. Thus for example if a 
little lithium chloride, LiCl, is mixed with some common salt 
the presence of the former can easily be detected with the spectro- 
scope. To do this the end of a platinum wire is dipped in the 
solution to be tested and then held in a non-luminous flame 
toward which the slit of the spectroscope is directed. One sees 
the lines of both lithium and sodium. 

787. Other Means of Examining Spectra. — Only a small 
number of elements give colored flames suitable for spectro- 
scopic study. By suitable means every element can be made to 
show a bright line spectrum. In general it is necessary for this 
purpose to heat the element to a higher temperature than that 
of the Bunsen flame. To do this we may make use of the electric 
arc; or we may cause sparks to pass between a platinum wire 
and a solution of the substance. In either case bright-line 
spectra are seen; but there is usually more or less difference 
between the flame, arc, and spark spectra of a given element. A 
still different method applicable to gases and volatile substances 
consists in rendering them luminous by the discharge of an 



506 Introduction to General Chemistry 

induction coil while they are contained in a glass tube under 
low pressure. The spectroscope is of great importance in the 
analysis of substances, and in several cases its use has led to the 
discovery of new elements. 

788. The Ether Wave Hypothesis. — Light is a vibratory 
disturbance in the so-called luminous ether, the waves being 
set up (according to hypothesis) by atomic vibrations. Light 
of a definite color is due to waves all of equal length. In so far 
as the visible spectrum is concerned, red light has the longest 
waves and violet light the shortest. Waves of intermediate 
lengths give all the intervening colors of the spectrum. An 
incandescent solid sends out waves of all lengths and gives 
therefore a continuous spectrum. On the other hand, a glowing 
gas emits waves of but very few wave-lengths, corresponding to 
the bright lines of its spectrum. The reason for this is thought 
to be that each distinct length of wave is caused by a single sort 
of vibrator. Very likely it is the electrons composing an atom 
(482) which act as the vibrators which set up the ether waves 
appearing to the eye as light. 

789. Dark-line Spectra. — If colored flame (say of sodium 
light) is interposed between a highly luminous solid and a spec- 
troscope one might expect to see a continuous spectrum with 
its yellow portion crossed by a still brighter yellow line. The 
effect, however, is quite the contrary: where we should expect 
the bright line of sodium a dark line appears instead. Under 
similar conditions the colored flame of any other element would 
in like fashion show a dark-line spectrum of that element, having 
a dark line corresponding to every bright line of its ordinary 
spectrum. The explanation of this curious fact seems to be the 
following: vibrators (atoms or electrons) which can emit light 
of a certain wave-length are also able to absorb light of this same 
wave-length. Therefore those wave-lengths of the light of the 
bright continuous spectrum which correspond to the bright lines 
of the glowing gas are absorbed in passing through the latter, 
so that in their places dark lines appear in the spectrum. 

790. The Composition of the Sun and the Stars. — The spec- 
trum of daylight (sunlight) shows a great number of dark lines. 



The Atmosphere and Related Topics 507 

These were discovered by Wollaston in 1802 and first mapped 
by Fraunhofer in 18 14; they are commonly called Fraunhofer 
lines. These lines correspond to the elements present in the 
sun's glowing atmosphere. Among the elements so indicated 
are many common on earth such as calcium, iron, hydrogen, and 
sodium. A few lines are present, however, which do not cor- 
respond to any known elements. Some of these are thought to 
be due to elements not as yet discovered on earth. 

We may even learn much of the chemical composition of the 
stars, which are in fact very distant suns, by a study of their 
dark-line spectra. 

791. The Discovery of Argon. — In the year 1890 there was 
probably not a single reputable chemist in the world who would 
have conceded the possibility of the existence in the air, to the 
extent of nearly 1 per cent, of a hitherto unknown gaseous 
element. The discovery of argon in 1894 by Rayleigh and 
Ramsay forms, therefore, one of the most interesting and sug- 
gestive chapters of modern chemistry. The whole story as 
related by Lord Rayleigh himself in the article on argon in the 
eleventh edition of the Encyclopaedia Britannica is well worth 
reading. Briefly stated the circumstances leading to the dis- 
covery are these: Lord Rayleigh, one of the world's greatest 
physicists, and Professor Ramsey (later Sir William), a brilliant 
chemist, both Englishmen, were engaged in a research on the 
determination of gas densities, in which results of the utmost 
accuracy were desired. Nitrogen was one of the gases studied. 
Now, as the student well knows, nitrogen can be prepared from 
various nitrogen compounds (513) as well as also (so it was 
thought) from the air by the removal of oxygen, carbon dioxide, 
and water vapor, and the minute amounts of other known gases. 
Rayleigh and Ramsey found that nitrogen extracted in this way 
from air had a density half of 1 per cent greater than nitrogen 
made from ammonia or other pure compounds of nitrogen. This 
difference in density was about fifty times as great as could be 
accounted for by experimental error! The only logical conclu- 
sion was that " atmospheric nitrogen" contained an appreciable 
proportion of a new gas denser than nitrogen and equally or 



508 Introduction to General Chemistry 

more inert chemically than the latter. A thoroughgoing search 
of the literature of nitrogen brought to light a most valuable 
clue. Nearly a century earlier Cavendish (292), the discoverer 
of hydrogen, had carried out an experiment in which atmospheric 
nitrogen mixed with an excess of oxygen was subjected to the 
prolonged action of electric sparks (566) . The oxides of nitrogen 
so formed were absorbed in alkali, and the volume of the nitrogen 
was thus reduced to 1/1 20 of that taken. But further diminution 
of volume could not be made to take place. 

Rayleigh and Ramsay's repetition of Cavendish's experiment 
with much more refined and larger apparatus enabled these 
scientists to obtain considerable amounts of the chemically 
inert residual gas. When this was freed from nitrogen and other 
known gases it was found to be a colorless, odorless gas 1.25 
times as dense as oxygen (1 liter weighs 1 . 78 g.). It was named 
argon and given the symbol A. 

Every effort to cause argon to combine with other elements 
or react in any way with any other substance proved completely 
futile. Furthermore argon could not be decomposed by any 
physical means. It was therefore classed as an element, but 
one without chemical affinity and therefore without chemical 
properties. 

792. The Molecular and Atomic Weights of Argon. — At 
standard conditions, 22.4 liters of argon weigh 39.9 g.; there- 
fore the molecular weight of the element is 39.9. Since argon 
does not combine with other t elements it is impossible, by 
methods so far discussed in this text, to determine its atomic 
weight. If, however, we can by some independent method find 
the number of atoms in a molecule of argon we have only to 
divide its molecular weight by this number to get its atomic 
weight. 

By means of calculations based on the kinetic-molecular 
hypothesis it has been found that if no expansion is permitted 
the amount of heat required to raise the temperature of one gram 
molecular weight of a gas one degree (the so-called molecular 
heat) should be three calories, if each molecule consists of a 
single atom. In this case the heat applied is wholly used 



The Atmosphere and Related Topics 509 

to increase the velocities of the molecules. If however the 
molecules are made up of two or more atoms each, then part 
of the heat is used up in increasing the internal energy of each 
molecule (that is, increasing the vibrational or rotational 
velocities of the atoms within each molecule with respect to 
one another) . Therefore the heat required to raise the tempera- 
ture of one gram molecular weight of a diatomic gas should 
theoretically be more than three calories. 

Now it has been found by experiment, in strict accord with 
theory, that the vapors of sodium and mercury, elements having 
single atom molecules (75), have molecular heats of three 
calories. No known gas or vapor has a smaller molecular heat 
than this value. On the other hand, just as theory predicts, 
oxygen, hydrogen, and nitrogen, all forming diatomic mole- 
cules, have molecular heats of about 4.9 calories. Gaseous 
elements and compounds with still more complex molecules show 
still larger molecular heats. When the molecular heat of argon 
was found to be three calories the conclusion was at once drawn 
that its molecules consist of single atoms and therefore that its 
atomic weight is identical with its molecular weight, namely 
39 .9. Indeed this conclusion is quite in harmony with the fact 
that argon does not unite with other elements, for it would be 
very strange if its atoms united with one another when they 
show no inclination to unite with other atoms. 

793. The Ratio of Molecular Heats of Gases.— The molec- 
ular heat referred to in the previous section is the so-called 
molecular heat at constant volume, C v . It represents the heat 
necessary to raise one gram molecular weight of the gas through 
one degree centigrade when the volume of the gas is kept con- 
stant. If the gas is allowed to expand at constant pressure 
while it is being heated, two more calories per molar weight are 
needed to raise the temperature one degree. These two calories 
are used up in the work of overcoming the external pressure 
during the expansion of the gas. The value so determined 
is the molecular heat at constant pressure, C p . Now it 
appears that the ratio of the two molecular heats C p C v for a 
given gas may be calculated from the velocity of sound in the 



510 Introduction to General Chemistry 

gas in question, and that the ratio is a relatively easy value to 
determine. As a result it is usual to find C p /C v instead of the 
separate values. For if C v for monatomic gases is equal to 
three, then the ratio of the molecular heats may be represented 
thus: 

C v 3 

The ratio of the molecular heats of a diatomic gas is in no case 
more than i .40. The ratio of the molecular heats is at its maxi- 
mum in the case of monatomic gases. 

794. The Story of the Discovery of Helium. — Of all the 
elements known in 1894 argon stood alone in refusing to unite 
chemically with any other element. Naturally this proved a 
matter of surpassing interest to chemists at the time. After 
the extensive laboratory experiments in which Rayleigh and 
Ramsay hoped to make compounds of argon had all failed, it 
occurred to these scientists that such compounds might possibly 
exist in nature. But if this were true, where should they seek 
these strange substances? Once more a search of old chemical 
literature furnished a clue. In a paper by the American chemist 
Hillebrand there appeared a statement that the mineral cleveite 
when heated gave off nitrogen. Before the discovery of argon 
it was considered sufficient identification of a gas as nitrogen 
to show that it was chemically inert and not one of the other 
known gases. Perhaps the "nitrogen" from cleveite might be 
argon! Ramsay repeated Hillebrand's experiment and obtained 
very readily an inert gas, but it was not nitrogen, since it would 
not unite with oxygen under the action of electric sparks (566). 
Moreover, it was not argon, since its density was but one-tenth 
as great as that of argon. It was plainly a new element, with 
a density only twice that of hydrogen; but like argon, entirely 
devoid of chemical affinity. The spectrum of the new gas showed 
a yellow line not far from that of sodium but not identical with 
that of any hitherto known element. Search of the literature 
of spectroscopy, however, revealed a most interesting fact. In 
1868 the astronomer Lockyer studied the photographs of spectra 
of the sun's corona (the incandescent atmosphere of the sun) 



The Atmosphere and Related Topics 511 

taken during a total eclipse and had noted the presence of a 
Fraunhofer dark line in the yellow of the spectrum which could 
not be identified as due to any known element. In explanation 
Lockyer suggested that this line was doubtless due to a new 
element to which he gave the name helium (Gr. helios, the sun). 
Careful comparison of the position of the dark line of the Lockyer 
photograph with the yellow line of the new gas showed beyond 
any doubt that the two were produced by the same element. 
Therefore the new gas was Lockyer's helium! Thus Rayleigh 
and Ramsay's search for a possible compound of argon led to the 
discovery on earth of an element whose existence in the sun had 
been announced thirty years earlier on the evidence of a photo- 
graph taken of an object 92 million miles away. 

795. Helium and the Alpha Rays of Radioactive Sub- 
stances. — We have already stated (480) that the alpha rays of 
radioactive substances are helium atoms. The further discus- 
sion of this scientifically important topic will be taken up in 
chapter xxxi. At this point we shall only add that minerals 
like cleveite, in which helium is contained, all contain radio- 
active substances, and that the helium present is beyond doubt 
the product of radioactive changes (867) . Helium is present in 
the atmosphere in minute amount, about four parts in a million. 

796. The Properties of Helium. — As we have stated, helium, 
like argon, is devoid of all chemical properties. It is a colorless, 
odorless gas of twice the density of hydrogen. One liter weighs 
0.18 g. and 22.4 liters weigh 4.0 g. Therefore its molecular 
weight is 4. Since the ratio of its two molecular heats (793) is 
1 .66, it is a monatomic gas and its atomic weight is therefore 4. 
Of all gases helium is the most difficult to liquefy. It has the 
lowest critical temperature, — 268 or 5 absolute. By cooling 
compressed helium by boiling liquid hydrogen and allowing it 
to expand in a liquerier similar in principle to that used for liquid 
air (776) it has been condensed to a liquid which boils at 4.3 
absolute. The boiling-point of helium under reduced pressure 
showed a temperature of about 3 absolute; this is up to the 
present time the lowest temperature ever produced experi- 
mentally. The liquefaction of helium was accomplished by 



512 Introduction to General Chemistry 

Dr. Kammerlingh Onnes, of the University of Leiden, in the 
year 1908. 

797. Helium Balloons. — The use of balloons in warfare is 
extremely precarious on account of the inflammability of the 
hydrogen with which they are filled. Helium being incom- 
bustible would be an ideal substitute for hydrogen if it could only 
be had in sufficient quantity. The lifting power of a helium 
balloon is eleven-twelfths of that of the same balloon filled with 
hydrogen. For a number of years it had been known that helium 
was often found in small proportions in natural gas, and when 
the matter of helium balloons began seriously to be considered 
in the last year of the war (19 18) it was found that a certain 
gas well in Texas produced gas containing nearly 1 per cent of 
helium. Calculation showed that this well also produced a 
sufficient total amount of helium to be of practical importance. 
The United States government appropriated $500,000 for a 
plant to produce pure helium, and at the end of the war pro- 
duction had already begun. It is not unlikely that immense 
dirigibles filled with helium will in the near future furnish a 
safe means of aerial travel. If this turns out to be the case it 
will be because two English scientists were keen to discover the 
cause of a seemingly trivial matter — the small discrepancy in 
the density of nitrogen as prepared in different ways! But this 
is the usual course of great practical discoveries, the scrupulous 
following up of every outstanding scientific inconsistency. 

798. Three Other Inactive Gases. — Shortly after the dis- 
covery of helium three other inert gases were discovered in 





TABLE XL 




Gas 


Symbol 


Atomic Weight 


Neon 


Ne 
Kr 
X 


20 


Krypton 

Xenon 


S3 
130 



minute amount in the atmosphere. Like helium and argon 
these new gases do not form any chemical compounds; they are 
also monatomic. These gases with their symbols and atomic 
weights are given in Table XL. 



The Atmosphere and Related Topics 513 

799. Summary and Conclusions. — The air may be considered 
as a great gaseous ocean into which flow many gases of terrestrial 
origin and from which are drawn the gases entering into the 
earth's chemical activity. It is evident that in addition to the 
several well-known components of the air there must also be 
present numerous other gases and vapors in minute proportions. 
The composition of the air remains nearly constant because of 
the approximate equality of formation and removal of the several 
components. In addition to the gases present, water vapor and 
dust are components of prime importance. Without dust (or gas- 
eous ions) we should have no clouds, and without clouds no rain, 
although we should have high humidity and tremendous dews. 
Viewed in the light of the matters considered in chapter xxviii 
we must see that the atmosphere is a disperse system consisting 
of a complex gaseous mixture filled with a suspension of minute 
solid particles (dust) and liquid droplets (water). The gaseous 
ions of the air become in part attached to these suspended 
particles giving to them plus and minus charges, corresponding 
to the charges on particles of colloidal solutions. 

The work on the liquefaction of gases so ably started by 
Faraday in 1823 was brilliantly completed when Kammerlingh 
Onnes liquefied helium in 1908. Every known gas has now been 
liquefied, and all save helium have been solidified. In the 
liquefaction of gases two principles are most noteworthy: first, 
the gas can exist as a liquid only below its critical temperature; 
second, a highly compressed and sufficiently cold gas will, upon 
being allowed to expand, fall greatly in temperature. 

The spectra of elements volatilized at high temperatures, as 
in a flame or by means of an electric spark or arc, appear as 
bright lines. Each element gives definite and characteristic 
lines. If white light from a glowing solid passes through an 
incandescent gas, the continuous spectrum of the former is seen 
to be crossed by dark lines which are the Fraunhofer spectrum 
of the gas. The spectroscope is of great importance for the 
detection of elements. By its aid the composition of the atmos- 
phere of the sun and of the stars has been definitely revealed. 



CHAPTER XXX 
SOME ADDITIONAL ELEMENTS AND THEIR COMPOUNDS 

800. Introduction. — The total number of elements known at 
the present time is 83 (if we exclude the products of radioactive 
change, chap, xxxii). These with their symbols and atomic 
weights are given in Table XLL In the foregoing chapters we 
have become acquainted with but little more than one-third of 
the elements. In the present chapter we shall study more or 
less briefly the chemistry of a number of additional elements. 
These, with those studied earlier, include all of the elements of 
practical importance. This list does not embrace several ele- 
ments which are fairly abundant but for which no technical uses 
have as yet been discovered. In chapter xxxi these remaining 
elements will also be considered. 

Our study of part of the elements has enabled us to develop 
and illustrate many principles and laws; but the student will 
naturally wonder whether the study of the remaining elements 
will bring forth an entirely new set of generalizations. We can 
hasten to assure him that this is not the case, for in the study 
of new elements and compounds he will encounter very few 
facts and phenomena that do not have their counterparts among 
those of more familiar substances. 

801. Boron. — Borax and boric acid (also known as boracic 
acid) are doubtless known to everyone. These are compounds 
of the non-metallic element boron. This element occurs plenti- 
fully as boric acid, H 3 B0 3 , and also as borates of sodium, magne- 
sium, and calcium, but never as free boron. It is found chiefly in 
Italy, California, and Thibet. Boric acid, which is volatile with 
steam, is present in the steam escaping from the earth in certain 
volcanic regions, especially in Italy. Sodium borate, or borax, 
which is easily soluble, occurs abundantly in the waters of some 
California and Nevada lakes. It is also found in the deposits 
left by the drying up of certain lakes of this same region. Borax 

514 



Some Additional Elements and Their Compounds 515 

is also made technically from the borates of magnesium and 
calcium. Boric acid and borax are the only compounds of this 



TABLE XLI 

International Atomic Weights, 191 7 



Element 



Aluminum . 
Antimony. . 

Argon 

Arsenic .... 
Barium. . . . 
Beryllium. . 
Bismuth . . . 

Boron 

Bromine . . . 
Cadmium. . 
Caesium. . . 
Calcium. . . 
Carbon .... 
Cerium .... 
Chlorine . . . 
Chromium . 

Cobalt 

Columbium 
Copper. . . . 
Dysprosium 
Erbium. . . . 
Europium. . 
Fluorine. . . 
Gadolinium 
Gallium. . . 
Germanium 

Gold 

Helium .... 
Holmium. . 
Hydrogen. . 
Indium. . . . 

Iodine 

Iridium. . . , 

Iron 

Krypton. . . 
Lanthanum 

Lead 

Lithium. . . 
Lutecium. . 
Magnesium 
Manganese . 
Mercury. . . 



Symbol 



Al 

Sb 

A 

As 

Ba 

Be 

Bi 

B 

Br 

Cd 

Cs 

Ca 

C 

Ce 

CI 

Cr 

Co 

Cb 

Cu 

Dy 

Er 

Eu 

F 

Gd 

Ga 

Ge 

Au 

He 

Ho 

H 

In 

I 

Ir 

Fe 

Kr 

La 

Pb 

Li 

Lu 

Mg 

Mn 

Hg 



Atomic 
Weight 



27.I 
I20. 2 
39-88 
74.96 

137-37 
9.1 

208.0 
II .O 
79.92 

1 1 2 . 40 

I32.8I 
4O.O7 
I 2 . OO5 

I40.2S 
35-46 
52.O 

58.97 

93- 1 
63-57 

162.5 

167.7 

152.0 
19.0 

157-3 
69.9 

72.5 
197.2 
4.00 

163.5 

1 .008 
114. 8 
126.92 
193- 1 
55 84 
82.92 
139.0 
207 . 20 
6.94 

1750 
24.32 

54-93 
200.6 



Element 



Molybdenum . . 
Neodymium. . . 

Neon 

Nickel 

Niton (radium 
emanation) . . 

Nitrogen 

Osmium 

Oxygen 

Palladium. . . . 
Phosphorus. . . 

Platinum 

Potassium 
Praseodymium 

Radium 

Rhodium 

Rubidium 

Ruthenium. . . . 
Samarium 

Scandium 

Selenium 

Silicon 

Silver 

Sodium 

Strontium. . . . 

Sulfur 

Tantalum 

Tellurium 

Terbium 

Thallium 

Thorium 

Thulium 

Tin 

Titanium 

Tungsten .... 

Uranium 

Vanadium. . . , 

Xenon 

Ytterbium . . . 

Yttrium 

Zinc 

Zirconium. . . , 



Symbol 



Mo 

Nd 

Ne 

Ni 

Nt 

N 

Os 

O 

Pd 

P 

Pt 

K 

Pr 

Ra 

Rh 

Rb 

Ru 

Sa 

Sc 

Se 

Si 

Ag 

Na 

Sr 

S 

Ta 

Te 

Tb 

Tl 

Th 

Tm 

Sn 

Ti 

W 

U 

V 

Xe 

Yb 

Yt 

Zn 

Zr 



Atomic 
Weight 



96.O 

1443 

20. 2 
58.68 

222.4 

I4.OI 
I9O.9 

16.OO 
I06.7 

3I-04 
195.2 

39.IO 
I4O.9 
226.0 
IO2.9 

85-45 
IOI.7 

I50.4 
44.I 
79.2 
28.3 

IO7.88 
23.OO 
87.63 
32.06 

l8l. 5 

127.5 
159-2 
204.O 

232.4 
168.5 
Il8. 7 

48.I 
184.O 
238.2 

5IO 
I30.2 

173-5 

S8.7 

65-37 
90.6 



element that are of any commercial importance; these are 
cheap substances and are produced in large quantities. 



516 Introduction to General Chemistry 

802. Boric Acid, H 3 B0 3 . — Boric acid forms white crystals 
which dissolve in 25 parts of cold or 3 parts of boiling water. 
It is a very weak acid, weaker even than carbonic acid. Its solu- 
tion is almost tasteless; it certainly has no sour taste. The 
solution is used extensively in medicine as an antiseptic lotion. 
A cold, saturated solution is an excellent eye wash. If alkali 
has by accident got into the eye, after the former has been washed 
out at once with water, boric acid solution should be copiously 
applied. In the past boric acid was rather extensively used as a 
preservative for foods, especially milk. Its use in this way is 
now prohibited, although it is not poisonous if taken internally 
in small amount. It is sometimes prescribed in medicine for 
internal use (dose 0.3 to 1 g.). 

Boric acid is readily decomposed by heat into water and 
metaboric acid, HB0 2 , tetraboric acid, H 2 B 4 7 , and finally 
boron trioxide, B 2 3 : 

H 3 B0 3 ->HB0 2 +H 2 0, 
4 HB0 2 ->H 2 B 4 7 +H 2 0, 

H 2 B 4 7 ->2B 2 3 +H 2 . 

The first two reactions take place at rather low temperatures, 
the last at a red heat. These reactions are analogous to those 
by which pyrophosphoric and metaphosphoric acids are formed 
from orthophosphoric acid (590). 

The ordinary salts of boric acid are those derived from meta- 
boric, or more commonly tetraboric, acid. Borax, sodium tetra- 
borate, Na 2 B 4 7 -ioH 2 0, forms colorless, glassy crystals. Boric 
acid is easily made from a solution of borax by adding hydro- 
chloric or sulfuric acid to a hot solution. On cooling, crystals of 
boric acid separate out: 

Na 2 B 4 7 +2HCl+5H 2 0->2NaCl+ 4 H 3 B0 3 . 

803. Borax. — Borax, or sodium tetraborate, Na 2 B 4 7 * ioH 2 0, 
is the only salt of boric acid of commercial importance. It is 
easily soluble in water, and its solution is alkaline by reason of 
hydrolysis (436). Borax is used extensively in the laundry for 
softening hard water (156). Hardness of water is due to the 



Some Additional Elements and Their Compounds 517 

presence of the bicarbonates and sulphates of calcium princi- 
pally and magnesium to a smaller extent. The bicarbonates 
cause temporary hardness, so called because they are decomposed 
and the carbonates precipitated when the water is boiled. 
The two sulphates cause what is called permanent hardness. 
Borax softens water by precipitating all the calcium and most of 
the magnesium present in hard water, as illustrated for the 
calcium salts by the following equations : 

Na 2 B 4 7 +Ca(HC03) 2 +5H 2 O^Na 2 C03+CaC03+4H 3 B03 

CaS0 4 +Na 2 C0 3 ->CaC0 3 +Na 2 S0 4 . (448) 

Upon being heated borax partially melts and at the same 
time gives off its water of hydration and in so doing swells to a 
spongy, viscous mass (intumesces) ; this finally melts com- 
pletely when all the water has been driven off, and it then 
resembles melted glass. The cold product, which is anhydrous 
borax, Na 2 B 4 7 , is a clear, colorless, brittle solid called borax 
glass. Beads made of borax glass are used in the same way as 
metaphosphate beads (596), which they closely resemble in 

TABLE XLII 

Colors of Borax Beads 



Element 


Oxidizing Flame 


Reducing Flame 


Copper 

Chromium 


Blue 

Green 

Yellow 

Brown 

Blue 

Amethyst 


Red particles of metal 
Green 


Iron 


Light green 


Nickel 


Gray particles of metal 

Blue 

Colorless 


Cobalt 

Manganese 







appearance. Various metallic oxides readily dissolve in fused 
borax glass, and in many cases the beads have colors char- 
acteristic of the added metallic element. If the element is 
readily oxidized or reduced the color of the bead will depend on 
whether it has been heated in an oxidizing or a reducing flame 
(783). Table XLII shows the colors of borax beads of several 
elements. 



518 Introduction to General Chemistry 

In explanation of the action of melted borax on oxides it 
may be supposed that reactions of the type illustrated by the 
following equation occur: 

Na 2 B 4 7 +CuO->2NaB0 2 +Cu(B0 2 ) 2 . 

Elements which give differently colored beads in the oxidizing 
and reducing flames do so because they undergo oxidation and 
reduction. Thus the yellow iron bead obtained in the oxidizing 
flame contains ferric metaborate, while the green one from the 
reducing flame contains a ferrous salt. The borates of copper 
and nickel are reduced to the corresponding metals in the 
reducing flame. It may be added that boron trioxide also forms 
a glass after being fused, and that it also readily unites with 
metallic oxides at high temperatures to form borates. Some 
useful optical glasses contain borates as essential ingredients. 
Borates are also used in making some kinds of enamels and 
glazes. In the welding of iron a little borax sprinkled on the 
hot metal dissolves the iron oxide always present and thus gives 
a clean surface for the weld. 

804. Silicon. — The element silicon is the second most abun- 
dant of the components of the earth's crust. It is estimated 
that about 50 per cent of the known part of the earth is oxygen 
and 25 per cent is silicon. This is not surprising when we realize 
that common quartz sand is essentially silicon dioxide, Si0 2 , and 
that numerous minerals composing the bulk of the earth are 
compounds of silicon. Among such minerals, besides quartz, 
we may mention clay, granite, agate, and opal as being well 
known. The element silicon does not occur free in nature, 
although it can be made by reduction of the oxide with carbon 
at the high temperature of the electric-arc furnace: 

Si0 2 +2C->Si+ 2 CO . 

Free silicon made in this way is only 90 to 98 per cent pure. It 
is a semimetallic solid somewhat resembling cast-iron in appear- 
ance and having a density of 2.35 and a melting-point of about 
1400 . It conducts electricity fairly well. It is not attacked 
by any of the common acids, with the exception of hydrofluoric, 



Some Additional Elements and Their Compounds 519 

with which it reacts to form silicon tetrafluoride, SiF 4 (270). 
Metallic silicon, as it is called by technical men, is too brittle 
to be of much importance for the manufacture of vessels or 
machines; but an alloy with iron containing about 14 per cent 
of silicon finds important use in chemical industry for the con- 
struction of vessels, pumps, etc., for the handling of strongly 
acid solutions. These alloys, known under the trade names of 
Duriron, Tantiron, etc., are extremely resistant to the action 
of acids. Ferro-silicon is an alloy of iron and silicon containing 
50 to 75 per cent of the latter and is used in making other silicon 
iron alloys containing less silicon. 

805. Compounds of Silicon with Hydrogen Carbon and 
Chlorine. — Silicon forms a great many compounds with hydro- 
gen, carbon, and chlorine, of much theoretical but little practical 
importance. Many of these compounds are analogous to certain 
compounds of carbon. Thus we have silico-methane, SiH 4 , a 
combustible gas, and silico-ethane, Si 2 H6, a spontaneously com- 
bustible liquid, corresponding to methane and ethane (643) 
respectively. Silicon tetrachloride, SiCl 4 , is a colorless liquid 
formed by the union of silicon and chlorine and also by the action 
of chlorine on a red-hot mixture of silica and carbon (charcoal) : 
Si0 2 +2C+2Cl 2 ^SiCl 4 +2CO . 

Silicon tetrachloride boils at 58 and resembles in many respects 
carbon tetrachloride, CC1 4 (644), of boiling-point 76 . Unlike 
the latter, the former reacts readily with water to form silicic 
acid, H 2 Si0 3 , and HC1: 

SiCl 4 +3H 2 0->H 2 Si0 3 + 4 HCl . 

The formation and properties of silicon tetrafluoride, SiF 4 , and 
of hydrofluosilicic acid and its salts have already been discussed 
(270-272). 

By the action of HC1 gas on heated silicon, silico-chloroform, 
SiHCl 3 , is obtained as a low-boiling colorless liquid. This is the 
analogue of chloroform, CHC1 3 (644), but unlike the latter it is 
readily decomposed by water. Silicon tetrachloride vapor 
reacts with heated silicon as follows: 

3SiCl 4 +Si->2Si 2 Cl 6 . 



520 



Introduction to General Chemistry 



The product, silicon hexachloride, is a liquid boiling at 148 , 
analogous to carbon hexachloride, C 2 C1 6 , boiling at 187 . There 
is little doubt that the structure of these substances is that 
represented by the following formulae: 



CI CI 

I I 

ci— c— c- 

I I 

CI CI 



-CI Cl- 



Cl CI 

I I 

-Si— Si- 

I I 
CI CI 



-CI 



Much more complicated silicon compounds are known, in 
some of which long chains (664) of silicon atoms occur. The 
resemblance between compounds of silicon and carbon extends 
to a great variety of compounds. By way of further illustration 
the following pairs of formulae of related substances are shown : 



CH 3 

I 
H 3 C— C— CH 3 

CH 3 
Tetramethyl-Methane 

OCOH 

I 
OCOH 

Oxalic Acid (665) 

OCOH 

I 
HOCOH 

I 
OCOH 

Mesoxalic Acid 



CH 3 

H 3 C-Si— CH 3 

CH 3 
Tetramethyl-Silico-M ethane 

OSiOH 

I 
OSiOH 

Silico-Oxalic Acid 

OSiOH 

I 
HOSiOH 

I 
OSiOH 

Silico-Mesoxalic Acid 



806. Silica or Silicon Dioxide, Si0 2 . — Silica is one of the 
most abundant of all minerals. It occurs in many forms both 
crystalline and amorphous. In a pure state it forms colorless, 
transparent, hexagonal crystals called quartz or rock crystal. 
Amethyst is a variety of quartz colored by manganese. The 
commonest kinds of sand are usually largely or wholly quartz. 
Sandstone is essentially quartz, red varieties being colored 



Some Additional Elements and Their Compounds 521 

with ferric oxide. A massive form of silica is called quartzite. 
Amorphous (581) varieties* constitute agate and onyx. Flint 
and opal are hydrated varieties of silica. 

The various forms of silica are extremely resistant to chemi- 
cal and physical changes by reason of the great hardness, high 
melting-point, and chemical inertness of this substance. These 
properties, together with the abundance and cheapness of various 
forms of silica, cause them to be used in enormous amounts for 
a great variety of purposes. As an abrasive silica finds use in 
grindstones, whetstones, sandpaper, and polishing material, the 
latter containing silica powder, tripoli, or infusorial earth (732). 
Mortar and concrete always contain quartz sand. Fire bricks for 
furnace linings are often made of silica. Various kinds of glass, 
including that made from pure quartz, will be considered later 
(808). 

807. Silicic Acid and Silicates. — Silica reacts with sodium 
carbonate (soda ash) at a red heat to form a viscous fluid con- 
sisting of sodium silicate, Na 2 Si0 3 . During the fusion carbon 
dioxide is evolved: 

Si0 2 +Na 2 C0 3 ->Na 2 Si0 3 +C0 2 . 

On cooling, sodium silicate solidifies to a glassy solid easily 
soluble in water. By adding hydrochloric acid to a solution of 
sodium silicate we get silicic acid, which under certain conditions 
appears as a gelatinous precipitate (a gel, 761). The com- 
position of this silicic acid is approximately represented by the 
formula H 4 Si0 4 . Its formation may be represented thus: 

Na 2 Si0 3 +H 2 0+2HCl->H 4 Si0 4 +2NaCl . 

Silicic acid is an extremely weak acid, but it will dissolve in 
sodium hydroxide solution to form a solution of sodium silicate. 
The solution is strongly alkaline by reason of hydrolysis. 

Silicic acid gel gradually loses water when heated, and in 
so doing it doubtless gives a series of derived acids and finally 
silica. In this respect it resembles phosphoric (590) and boric 
acids (802). Theoretically we have the following possibilities 



522 



Introduction to General Chemistry 



as one or two molecules of H 4 Si0 4 lose water. (In the following 
expressions the water lost is omitted for the sake of brevity.) 

H 4 Si0 4 ->H 2 Si0 3 ^Si0 2 
2H 4 Si0 4 ->H 6 Si 2 7 ^H 4 Si 2 6 ->H 2 Si 2 5 ->2Si0 2 . 

Still more complex acids would result from the dehydration 
of three or more molecules of silicic acid. Although none of 
these acids has been made in pure form, salts of several of them 
are found as minerals. Silicon dioxide is of course the anhydride 
(3*3) of all the possible derivatives of silicic acid, and the formulae 
of all the latter may be written in the general form (H 2 0) w (Si0 2 ) m , 
where n and m are integers. The formulae of the salts may also 
be written similarly, as illustrated in Table XLIII, which shows 
the composition of several minerals and the acids from which 
they are derived. 

TABLE XLIII 



H 4 Si0 4 . 
H 2 Si0 3 . 
H 4 Si0 4 . 
H 2 Si0 3 . 

H 4 Si 3 8 



Zn 2 Si0 4 
CaSi0 3 

2 H 2 KAl 3 (Si0 4 ) 3 

H 2 Mg 3 (Si0 3 ) 4 

2KAlSi 3 8 



2ZnO-Si0 2 

CaOSi0 2 

K 2 0- 3 Al 2 3 -6Si0 2 -2H 2 

3 MgO- 4 Si0 2 -H 2 

K 2 OAl 2 3 -6Si0 2 



Willemite 

Wollastonite 

Mica 

Talc 

Orthoclase 



808. Glass. — There are a great many kinds of glass, nearly 
all of which are silicates, although borates are also common 
ingredients. The basic constituent elements include sodium 
potassium, calcium, aluminum, zinc, lead, and other metals less 
commonly used. Water glass having approximately the com- 
position Na 2 Si 4 9 is made by dissolving infusorial earth (732) in 
hot sodium hydroxide solution. It comes on the market as a 
heavy syrup which gives on evaporation a glasslike solid easily 
soluble in water. It is used as a component of laundry soaps, 
as a cement, as an oil-proof glaze for lard barrels, and for a 
number of other purposes. 

Glass made from the silicates of sodium and calcium is used 
for window panes, bottles, etc. The ordinary soft glass tubing 
of the laboratory is also made of this kind of glass. Hard glass, 
which contains potassium in place of the sodium in soft glass, 



Some Additional Elements and Their Compounds 523 

fuses at a higher temperature than the latter. Flint glass, a 
potassium-lead silicate, has a high index of refraction and on 
this account possesses great brilliancy. It is used for cut-glass 
ware and lenses. 

Glassware for chemical use, beakers, flasks, etc., should 
possess several special properties. It must be practically 
insoluble in water, acids, and alkalies, must not crack when 
heated, and must be thick enough to withstand ordinary handling 
without breaking. The cracking of glass when heated is the 
result of rapid local expansion. The thinner the glass the less 
liable it is to crack when heated; but if ordinary glass is made 
thin enough to withstand heating in the usual way it is very 
fragile. Before the war nearly all chemical glassware came from 
Germany and Bohemia. There were two principal kinds, typi- 
cally represented by Kavalier glass and Jena glass, the latter 
being much better in quality than the former. The composition 
of each of these is shown in Table XLIV. 

TABLE XLIV 

Percentage Composition of Glass for Laboratory Ware* 



Kavalier 



Jena 



Pyrex 



Nonsol 



A1 2 3 - 

Fe 2 3 

ZnO. 

CaO. 

MgO. 

K 2 0. 

Na 2 0. 

Si0 2 . 

BA- 

As A 



o. 14 



8.70 



7.90 

7 . 10 

75 -9° 



4. 20 

0.25 

10.90 

0.63 

o. 21 

o.37 

7 50 

64.70 

10.90 

o. 14 



2 .00 
0.25 



o. 29 
0.06 
o. 20 

4.40 
80.50 
11.80 

o. 70 



2 


50 


O 


23 


7 


80 





79 


3 


40 





30 


10 


90 


67 


30 


6 


20 



* Walker & Smith, U.S. Bureau of Standards, Bulletin 107, 1018. 



During the war several American manufacturers succeeded 
in producing first-class chemical glass. The composition of two 
sorts, Pyrex and Nonsol, is shown in Table XLIV. Pyrex glass 
has so small a coefficient of expansion that beakers and flasks 
can be made with thick walls and still be far less liable to crack 
with change of temperature than the best German glass. Several 



524 Introduction to General Chemistry 

other sorts of American chemical glass are also of excellent 
quality, so that today we have better domestic ware than was 
heretofore imported. 

Quartz glass is made from pure silicon dioxide (806) melted 
in an electric furnace. This glass resembles ordinary glass in 
appearance but differs from the latter in having an extremely 
high melting-point and an extremely small coefficient of expan- 
sion. White-hot quartz will not crack when plunged into water. 
This glass finds important uses in both scientific and technical 
chemistry. 

809. Tin and Its Compounds. — The metal tin (Sn=ng) was 
known from very early times, its Latin name being stannum. 
It is widely used in making a number of alloys, such as bronze 
and bell metal (tin and copper), solder (tin and lead), and tin 
amalgam (tin and mercury used in making mirror backs) . Sheet 
iron coated with tin and known as tin plate is made in immense 
quantities for the manufacture of common tinware. Tin in the 
form of thin sheets is tin foil, but the common grade of this 
article is usually largely lead. 

Tin is a rather soft metal of low melting-point, 23 2 °. It is 
permanent in air and in water, and, as shown by its position in 
the displacement series (Table XIX, 492), it is a rather inactive 
element. It dissolves somewhat slowly in hydrochloric acid, 
forming stannous chloride, SnCl 2 , 

Sn+ 2 HCl->SnCl 2 +H 2 . 

The hydrate SnCl 2 • 2H 2 0, easily soluble in water, is known in 
commerce as tin salt and is used in dyeing. It is frequently used 
in the laboratory as a reducing agent. Stannous sulfate, SnS0 4 , 
and nitrate, Sn(N0 3 ) 2 , can also be made. The action of chlorine 
gas on melted tin gives stannic chloride, SnCl 4 , 

Sn+2Cl 2 ->SnCl 4 . 

The product is a colorless liquid boiling at 114 , which forms 
with water the crystalline hydrates SnCl 4 • 5H 2 and SnCl 4 • 8H 2 0. 
These salts are also used as mordants in dyeing, a mordant being 
a substance which by combining with the fiber fixes the dye upon 
the latter. Numerous other stannic salts are known. 



Some Additional Elements and Their Compounds 525 

Thus it will be seen that tin forms two series of salts: the 
stannous, in which the metal is bivalent, and the stannic, in 
which it is quadrivalent. Stannous salts form stannous ions, 
Sn ++ , which by oxidation give stannic ions, Sn 4+ . The latter 
by reduction give the former. Stannous solutions are strong 
reducing agents. 

Tin forms two oxides, stannous, SnO, and stannic, Sn0 2 . 
The corresponding hydroxides are precipitated when a stannous 
or a stannic solution is treated with the equivalent amount of 
sodium hydroxide. Each precipitate is soluble in an excess of 
the alkali, thus showing that the hydroxides of tin have acidic 
as well as basic properties. Like aluminum hydroxide they are 
amphoteric (177). Stannous hydroxide, Sn(0H) 2 , yields sodium 
stannite, Na 2 Sn0 2 ; while stannic hydroxide, also called stannic 
acid, gives sodium stannate, Na 2 Sn0 3 . A solution of sodium 
stannite is a very powerful reducing agent, since it is readily 
oxidized to sodium stannate. An isomer (647) of stannic acid, 
called metastannic acid, is formed by the action of hot nitric 
acid on tin. This is a white solid, insoluble in water and ni- 
tric acid, and differs greatly from the other form. Metastannic 
acid forms complex salts such as Na 2 Sn 5 0n*4H 2 0. Hydrogen 
sulfide precipitates brown stannous sulfide from solutions of 
stannous chloride and yellow stannic sulfide from solutions of 
stannic chloride. Both precipitates are soluble in concentrated 
hydrochloric acid. Stannic sulfide is dissolved by ammonium 
sulfide, owing to the formation of ammonium sulfo-stannate. 
Stannous sulfide does not dissolve appreciably in pure ammonium 
sulfide but does so in yellow ammonium sulfide (607), forming 
also ammonium sulfo-stannate: 

SnS 2 +(NH 4 ) 2 S ->(NH 4 ) 2 SnS 3 , 
SnS '+(NH 4 ) 2 S 2 ^>(NH 4 ) 2 SnS 3 . 

810. Arsenic. — In its chemistry arsenic (As = 75) is much 
more like the non-metallic than the metallic elements. In 
many of its compounds it bears a close resemblance to phos- 
phorus. Arsenic is an abundant element, occurring as a mineral 
both free and in combination with oxygen and sulfur, and with 



526 Introduction to General Chemistry 

iron, copper, and many other metals. The trioxide As 2 3 , 
commonly called white arsenic, is obtained in large amounts as 
a by-product of the smelting of ores of copper and some other 
metals. 

The free element is easily prepared by heating the trioxide 
with carbon; the vapor of arsenic so formed condenses on cooling 
to nearly black crystals. The latter are semimetallic in appear- 
ance and, like metals, conduct electricity well. Arsenic is not 
oxidized by air at ordinary temperatures but burns when strongly 
heated, forming the trioxide. When heated in a stream of 
chlorine, arsenic forms a trichloride, AsCl 3 , a colorless liquid 
boiling at 130 . This hydrolyzes almost completely when 
treated with much water, thus showing that the corresponding 
hydroxide, As(OH) 3 , is a very weak base. 

On the other hand the hydroxide and also the trioxide dis- 
solve readily in dilute alkalies to form salts in which As(OH) 3 
plays the part of an acid, called arsenious acid. Two series of 
salts are known: one derived from H 3 As0 3 , the other from 
HAs0 2 . Arsenious acid, like phosphorous acid, H 3 P0 3 (588), 
is a good reducing agent. It is Oxidized by halogens or nitric 
acid to form arsenic acid, H 3 As0 4 : 

H 3 As0 3 +H 2 0+Cl 2 ->H 3 As0 4 +2HCl . 

Arsenic acid, which is derived from the pentoxide As 2 5 , is a 
very soluble, moderately strong acid closely resembling phos- 
phoric acid (590). Like the latter, it loses water when heated, 
giving first pyroarsenic acid, H 4 As 2 7 , and then metarsenic acid, 
HAs0 3 . Arsenic acid resembles phosphoric acid in giving with 
ammonium molybdate solution a yellow precipitate of am- 
monium arseno-molybdate, insoluble in nitric acid (597). 

The numerous salts of arsenic acid xlosely resemble the 
corresponding salts of phosphoric acid. In fact Mitscher- 
lich in the year 18 19 discovered that corresponding salts, e.g., 
H 2 KP0 4 and H 2 KAs0 4 , were identical in crystalline form. Such 
compounds were said to be isomorphous. Later Mitscherlich 
declared that if two substances are isomorphous they must have 
the same numbers of atoms similarly arranged in the molecule. 






Some Additional Elements and Their Compounds 527 

•This principle enabled him to discover the correct formulae of 
many compounds of little-known elements at a time when much 
uncertainty prevailed regarding valence and atomic weights. 

The action of hydrogen sulfide on a solution of arsenious 
acid acidified with HC1 gives a precipitate of yellow arsenic 
trisulfide, As 2 S 3 (608). Acidified solutions of arsenic acid or 
its salts give under similar conditions arsenic pentasulfide, 
As 2 S s , along with some trisulfide and sulfur, according to experi- 
mental conditions. These sulfides are insoluble in water and 
hydrochloric acid but are easily soluble in concentrated nitric 
acid, which changes the substances into sulfuric and arsenic 
acids. The sulfides also dissolve in ammonium sulfide to form 
soluble sulfarsenite, (NH 4 ) 2 AsS 3 , and sulfarsenate, (NH 4 ) 2 AsS 4 . 

Arsenic and hydrogen form a gaseous compound called 
arsine, AsH 3 . This is produced along with hydrogen when any 
soluble arsenic compound is added to a mixture of zinc and 
hydrochloric acid. It is an extremely poisonous gas and must 
be handled with great caution. • Arsine is readily decomposed 
into its elements by heat, so that if it is passed through a glass 
tube heated locally a dark mirror-like deposit of arsenic is 
formed in the tube. Marsh's test for arsenic is based on these 
reactions. 

Arsenic compounds are -all poisonous; one-fifth of a gram 
of the trioxide usually proves fatal. A good antidote consists 
of freshly prepared ferric hydroxide, which combines readily 
with arsenious acid to form an insoluble and therefore little- 
poisonous compound. 

811. Antimony. —Antimony (Sb= 120) is classed as a metallic 
element because of the decidedly metallic properties of the free 
element and its ability to form salts with acids such as sul- 
furic and nitric. Nevertheless its hydroxides are also acidic 
and give rise to salts with bases, just as do the hydroxides of 
arsenic. Antimony is a widely distributed and rather common 
element occurring usually in combination with oxygen, sulfur, 
or metallic elements. China is the largest single source of the 
element. The metal forms brittle, silver-white crystals, melting 
at 630 . It forms useful alloys with many other metals. Type 



528 Introduction to General Chemistry 

metal, an alloy with lead and bismuth, expands at the moment 
of solidifying in the mold and thus produces type with sharp, 
clear-cut edges. Britannia, made from tin, antimony, and copper, 
is used as a cheap substitute for silver plate. Babbitt metal, 
used for antifriction bearings, consists of antimony and tin with 
some lead and copper. 

Antimony unites readily with chlorine to form the trichloride 
SbCl 3 , white crystals melting at 73 ° and boiling at 223 . The 
chloride is to be considered the salt of the weak base Sb(OH) 3 ; 
it is soluble in aqueous hydrochloric acid but is hydrolyzed by 
pure water. With a small proportion of water a white precipitate 
of sparingly soluble antimony oxychloride, SbOCl, is formed, 

SbCl 3 +H 2 O^SbOCl+2HCl . 

This reaction is reversible, SbOCl dissolving readily in concen- 
trated HC1. The univalent radical SbO present in several salts 
is called antimonyl; therefore SbOCl is also called antimonyl 
chloride. 

Antimony nitrate, Sb(N0 3 ) 3 , and antimony sulfate, Sb 2 (S0 4 ) 3 , 
are typical salts as far as their composition is concerned, but they 
are more or less completely hydrolyzed by water, depending 
on the proportion and temperature of the latter. The double 
salts potassium antimonyl tartrate, K(SbO)C 4 H 4 6 (see 665), 
and sodium antimony fluoride, NaSbF 4 , are easily soluble salts 
which are not hydrolytically decomposed by water. Both of 
these salts are important mordants (809), and the former, com- 
monly known as tartar emetic, is also used in medicine. In 
both cases stable complex ions (538) are formed; these are 
(SbO)C 4 H 4 6 ~ and SbF 4 ~ respectively. 

Antimony forms three oxides: Sb 2 3 , Sb 2 4 , and Sb 2 O s . 
The first and second can be made by direct union of antimony 
with oxygen, the second being the one stable in air at a red heat. 
The third, made indirectly, loses oxygen when strongly heated, 
giving Sb 2 4 . The trioxide corresponds to a hydroxide, Sb(OH) 3 , 
or H 3 Sb0 3 , which by reason of its faintly acidic nature is known 
as antimonous acid. The latter unites with bases to form salts 
such as Na 3 Sb0 3 and NaSb0 2 • 3H 2 0. Antimony pentoxide is 



Some Additional Elements and Their Compounds 529 

the anhydride of antimonic acid, H 3 Sb0 4 , a white powder spar- 
ingly soluble in water. This weak acid loses water when heated, 
giving in turn pyroantimonic acid, H 4 Sb 2 7 , metantimonic acid, 
HSb0 3 , and finally Sb 2 5 . Potassium hydrogen pyroantimonate, 
K 2 H 2 Sb 2 7 , is one of the best-known salts of these acids. Its 
solution gives with sodium salts a precipitate of sodium hydrogen 
pyroantimonate, Na 2 H 2 Sb 2 7 • 6H 2 0. 

Acid solutions of antimony salts give with hydrogen sulfide 
brick-red precipitates of antimony trisulride, Sb 2 S 3 , or penta- 
sulfide, Sb 2 S 5 . These sulfides are almost insoluble in water and 
cold dilute acids but dissolve easily in potassium hydroxide by 
reason of reactions like the following: 

2Sb 2 S 3 + 4 KOH->3KSbS 2 +KSb0 2 +2H 2 . 

The soluble salt KSbS 2 is called potassium sulfantimonite. 

Antimony trisulfide also dissolves in yellow ammonium 
sulfide to form soluble ammonium sulfantimonate, (NH 4 ) 3 SbS 4 . 

Antimony forms a gaseous hydride, SbH 3 , called stibine. 
This gas resembles arsine closely and is also decomposed by 
heat, with the formation of a black deposit of free antimony. 

812. Bismuth. — Bismuth (Bi=2o8) is strictly metallic in 
properties, in distinction from arsenic, which is non-metallic, 
and antimony, which stands midway between the two in this 
respect. In its ordinary compounds bismuth is trivalent, but 
products probably pentavalent also exist. The element is much 
scarcer than either arsenic or antimony and commands a much 
higher price. It occurs free (native) and as oxide or sulfide, 
often admixed with ores of copper or lead. Bolivia is the 
chief source of bismuth. A considerable amount of it is obtained 
in the United States as a by-product in the electrolytic refining 
of lead. The metal forms rather hard, brittle crystals, silver- 
white of a reddish tint. The melting-point is 269 . The metal 
is used in the making of easily fusible alloys. Thus Rose's metal, 
consisting of tin 1, lead 1, and bismuth 2 parts, melts at 94 ; 
Wood's metal, tin 1, lead 2, bismuth 4, and cadmium i, melts 
at 61 °. Fusible alloys are used for fire-sprinkler nozzles, electric 
fuse wires, etc. 



530 Introduction to General Chemistry 

Bismuth forms with acids salts in which this element acts 
as a trivalent positive ion. Bismuth trichloride, BiCl 3 , melting 
at 2 2 7 and boiling at 428 , is made by the direct union of chlorine 
with the metal, or by the action of aqua regia (562) on the lat- 
ter. The salt is soluble in water, but when diluted the salt 
hydrolyzes and forms a white precipitate of bismuth oxychloride, 
BiOCl (bismuthyl chloride or basic bismuth chloride). Bismuth 
nitrate, Bi(N0 3 ) 3 , results from the action of nitric acid on the 
metal. It is soluble in water or dilute nitric acid but is hydro- 
lyzed by much water, forming the basic nitrate BiON0 3 (also 
called bismuth subnitrate). This product is a nearly insoluble 
white powder used both internally and externally in medicine. 
Quite large doses are easily tolerated. Bismuth sulfide, Bi 2 S 3 , 
is formed as a brown precipitate when acid solutions of bismuth 
salts are treated with hydrogen sulfide. Sodium hydroxide or 
ammonia precipitates white bismuth hydroxide, Bi(OH) 3 , from 
a bismuth salt solution. Bismuth does not form a hydride. 

813. Molybdenum. — Molybdenum (Mo = 96) is found in 
nature chiefly as molybdenite, MoS 2 , a mineral closely resembling 
graphite (630) in appearance. When the sulfide is heated in 
air (roasted) it is oxidized to the trioxide Mo0 3 . The latter 
heated with hydrogen or carbon gives the metallic element. 
Molybdenum is a hard but malleable metal of very high melting- 
point (over 2400 ) . It is not oxidized by air at ordinary tempera- 
tures. The pure metal is not commercially important, but an 
alloy with iron, ferro-molybdenum, is made on a rather large 
scale. This alloy finds use in the making of certain kinds of high- 
grade tool steel. 

The trioxide Mo0 3 is the most important of the several 
oxides. It is the anhydride of molybdic acid, H 2 Mo0 4 • H 2 0, 
which consists of yellow crystals difficultly soluble in water. The 
hydrate loses water when heated gently, giving white H 2 Mo0 4 . 
Sodium molybdate, Na 2 Mo0 4 * ioH 2 0, resembles Glauber's salt, 
Na 2 S0 4 - ioH 2 0, in appearance. Molybdic acid resembles silicic 
acid in forming salts of complex formulae (807). Thus for 
example a sodium salt, Na 6 Mo 7 24 • 2 2H 2 0, and an ammonium 
salt, (NH 4 ) 6 Mo 7 24 *4H 2 0, are known. The latter salt dis- 



Some Additional Elements and Their Compounds 531 

solved in an excess of nitric acid gives the so-called ammonium 
molybdate solution, extensively used in the detection and quanti- 
tative analysis of phosphates (597). With this reagent a nitric 
acid solution of a phosphate or phosphoric acid gives quantita- 
tively a yellow precipitate of ammonium phospho-molybdate, 
(NH 4 ) 3 P0 4 -iiMo(V6H 2 0. Arsenic acid solutions give the 
corresponding arsenic compound under similar conditions (810). 

814. Tungsten. — The element tungsten (W=i84) is of con- 
siderable practical importance. In the United States it is found 
chiefly in Colorado. Its principal ore is Wolframite, a tungstate 
of iron and manganese. The separate tungstates FeW0 4 and 
MnW0 4 also occur as minerals. .Metallic tungsten is made in 
very pure form by the reduction of its oxide, W0 3 , by hydrogen 
at a high temperature. By reason of the extremely high 
melting-point of the metal the product is obtained as a powder. 
To get this into compact malleable and ductile form the powder 
is made into bars by powerful compression; these are then again 
heated electrically and the granules caused to melt together. 
Pure metallic tungsten, first made by Cooledge in 191 1, is a 
hard, silver-white, and very ductile metal. Its most important 
use is for the manufacture of the filaments of incandescent 
electric lamps. Its great value for this purpose depends on the 
fact that its melting-point, 3 267 °, is higher than that of any other 
metal. It may be added that the efficiency of an electric lamp 
(ratio of light energy to electrical energy) increases greatly with 
the temperature to which the filament can be heated. 

Ferro-tungsten (70 per cent W, 30 per cent Fe), made on a 
large scale by an electric-furnace process, is used in the manu- 
facture of special steels. Tungsten forms two oxides, W0 2 and 
W0 3 , the latter being far more important. The trioxide is the 
anhydride of tungstic acid, H 2 W0 4 , a yellow powder insoluble in 
water. This acid forms simple tungstates such as Na,W0 4 • 2ITO, 
K 2 W0 4 , CaW0 4 , PbW0 4 , and the tungstates of iron and man- 
ganese already mentioned. Calcium tungstate, CaW0 4 , gives 
a fluorescent light with X-rays and finds important use for the 
manufacture of fluoroscopic screens. Sodium tungstate is used 
to render cotton cloth slow-burning ("fireproof"). 



532 Introduction to General Chemistry 

815. Selenium and Tellurium. — The elements selenium 
(Se=79) and tellurium (Te=i27.5) are more closely related to 
sulfur than to any others. The first, selenium, is essentially 
non-metallic in its behavior. The second, tellurium, in free 
form is a silver-white crystalline metal which, in addition to 
forming some salts in which it is the basic ion, also forms two 
well-characterized acids. These elements are not very common. 
Practically they are obtained as by-products in the refining of cop- 
per. No important use has been found for tellurium. Selenium 
has two uses, one in the manufacture of red glass, the other in the 
construction of light-sensitive electrical apparatus. The electri- 
cal conductivity of selenium varies greatly with the intensity of 
its illumination, so that by ingenious use of this principle pic- 
tures, etc., have been transmitted long distances by wire. Lights 
on beacons and buoys may be turned on at dark by the use of 
selenium cells. Selenium and tellurium form compounds analo- 
gous to H 2 S by methods similar to those that yield the latter gas. 

Hydrogen selenide, H 2 Se, is an ill-smelling gas, dangerously 

irritating to mucous membranes and temporarily paralyzing 

the sense of smell. Hydrogen telluride, H 2 Te, is a similar gas. 

Selenium and tellurium both burn in air or oxygen when ignited 

and form white solid dioxides, Se0 2 and Te0 2 . These oxides 

both unite with water to form selenous acid and tellurous acid 

respectivelv : 

Se0 2 +H 2 0->H 2 Se0 3 , 

Te0 2 +H 2 0->H 2 Te0 3 . 

These acids are the analogues of sulfurous acid, H 2 S0 3 (611), 
but are more stable than the latter in that both are white 
crystalline solids. They form typical salts with bases; not 
only normal salts like K 2 Se0 3 and K 2 Te0 3 , but also acid salts 
like KHSe0 3 and KHTe0 3 . 

Unlike sulfurous acid, these acids are oxidizing agents, since 
they are reduced to selenium and tellurium respectively by 
reducing agents, as illustrated in the following equation: 

H 2 Se0 3 +2S0 2 +H 2 0->Se+ 2 H 2 S0 4 . 

On the other hand, powerful oxidizing agents convert selenous 
and tellurous acids into selenic acid, H 2 Se0 4 , and telluric acid, 



Some Additional Elements and Their Compounds 533 

H 2 Te0 4 , respectively. These acids form white crystals and are 
analogous in composition to sulfuric acid. Like the latter, they 
form both neutral and acid salts, e.g., potassium acid tellurate, 
KHTe0 4 . The selenates and tellurates are oxidizing agents 
tending to pass into selenites and tellurites. Thus, for example ? 
potassium tellurate, when strongly heated, gives the tellurite 
and oxygen, or when 'heated with concentrated hydrochloric 
acid it gives potassium tellurite and chlorine. Telluric acid is 
decomposed by heat into its anhydride, tellurium trioxide, Te0 3 , 
and water. 

816. Cobalt and Nickel. — The metals cobalt (00 = 58.97) 
and nickel (Ni=58.68) resemble iron in both their physical 
and their chemical properties but are less active chemically, 
corresponding to positions below iron in the displacement series 
(492). Nickel is by far the more important of the two com- 
mercially, although cobalt commands the higher price. The 
two metals usually occur together, nickel being the more abun- 
dant. They are commonly found with iron, copper, or silver as 
sulfides or arsenides. Large deposits of nickel and cobalt occur 
in the province of Ontario, Canada. 

A unique method of extracting nickel from its ores is based 
on the fact that carbon monoxide, CO (632), unites with finely 
divided metallic nickel at 8o° to form a compound, Ni(CO) 4 , 
nickel carbonyl, which is gaseous above 44 . The ores, after they 
have been roasted to remove sulfur and arsenic and to convert 
the metals into oxides, are first extracted with sulfuric acid to 
remove copper and are then reduced to metal by water-gas, a 
mixture of H 2 and CO, at 400 . The residue, cooled to 8o°, is 
exposed to a stream of highly compressed CO, and the volatile 
Ni(CO) 4 which passes off is, by heating it to 200 , decomposed 
into pure metallic nickel and carbon monoxide, 

Ni(CO) 4 ^Ni+ 4 CO . 

This is the Mond process. Other processes of refining nickel are 
also in use. 

Nickel is so extensively used for plating copper, brass, and 
iron that its appearance and permanence in the air are known to 
everyone. It is harder than iron and melts at 1450 . The metal 



534 Introduction to General Chemistry 

finds one of its most important uses in the manufacture of 
nickel steel, of which it may constitute several per cent. Cobalt 
may also be present to advantage in such steel. Nickel steel is 
very hard and tough and is largely used for armor plate. 

Our five-cent coins, "nickels," are an alloy of 25 per cent 
nickel and 75 per cent copper. The well-known properties 
of this alloy need no comment. The alloy called German 
silver consists of nickel, copper, and zinc. Monel metal is 
Ni 68 per cent, Cu 30 . 5 per cent, and Fe 1.5 per cent. It 
resembles the alloy of our nickel coins but is harder and more 
resistant to chemicals. It is obtained by direct smelting of 
copper-nickel-iron ores. 

Cobalt resembles nickel in appearance and properties. It is 
not much used as a metal. Its oxide, CoO, is used in making 
blue glass, and since no other means is known of making a good 
blue glass this element is of considerable importance. Blue 
glazes for porcelain and a blue pigment, smalt, are also made 
from cobalt oxide. 

817. Salts of Cobalt and Nickel. — The bivalent oxides and 
hydroxides of cobalt and nickel are basic and have no acidic 
properties. They form salts with practically all acids. In its 
salts nickel is always and cobalt usually bivalent, although a 
few trivalent salts of the latter element are known. The salts 
of these elements correspond closely in composition to those of 
ferrous iron. Of the commoner salts of (bivalent) cobalt and 
nickel we may mention the fluorides, chlorides, bromides, iodides, 
nitrates, sulfates, carbonates, and phosphates. These salts, 
with the exception of the carbonates and phosphates, are readily 
soluble in water and as a rule form crystals with six molecules of 
water to one of the salt, e.g., Co(N0 3 ) 2 • 6H 2 and NiCl 2 • 6H 2 0. 
Nickel sulfate forms besides the hexahydrate a heptahydrate, 
NiS0 4 -7H 2 0, isomorphous (810) with MgS0 4 • 7H 2 (144). 
Cobalt sulfate also forms a hexahydrate and a heptahydrate 
the latter isomorphous with FeS0 4 • 7H 2 (173). Many 
double salts (175) are also known. Ammonium nickel sul- 
fate, (NH 4 ) 2 S0 4 • NiS0 4 • 6H 2 0, is used in electroplating. 

The hydrated salts of cobalt are red in crystalline form and 
also in solution, in which latter Co ++ ions are present. Many 



Some Additional Elements and Their Compounds 535 

anhydrous cobalt salts are blue. If one uses a dilute cobalt 
chloride solution, pale red in color, as ink, the writing will be 
practically invisible when air dry. If the paper is heated the 
writing appears in easily legible blue characters. The effect of 
the heating is to drive off the water of hydration. A writing- 
fluid of this kind is known as sympathetic ink. Cobalt forms a 
large number of compounds called cobalt amines. These are 
complex substances containing ammonia and trivalent cobalt. 
Hexamine cobaltic chloride, (NH 3 ) 6 CoCl 3 , formed by the 
action of ammonia and oxidizing agents on a solution of cobaltous 
chloride, CoCl 2 , is a typical example. Solutions of cobalt 
amines must contain complex ions (538), since they do not give 
the ordinary reactions of solutions of simple cobalt salts. Nickel 
does not form analogous compounds. 

Hydrated nickel salts are green and form green solutions 
in which bivalent Ni ++ ions are present. Simple cobalt and 
nickel solutions give precipitates with alkaline hydroxides, 
carbonates, phosphates, and sulfides. These are all soluble in 
ammonia solution by reason of the formation of complex ions 
of various sorts. 

818. Platinum. — Platinum (Pt=i'95.'2) occurs only in the 
metallic state, often as an alloy with the rarer metals osmium 
and iridium. It comes chiefly from the Ural Mountains in 
Russia. Platinum is almost indispensable in chemical labora- 
tories, where it is used as crucibles, dishes, tubes, etc. Its 
value for such uses arises from the fact that it is very inactive 
and has a high melting-point. It is not attacked by any of the 
common acids singly, although it is slowly converted into its 
chloride, PtCl 4 , by aqua regia (562), a mixture of hydrochloric 
and nitric acids. Platinum is an active catalytic agent for many 
reactions, such as the union of hydrogen and oxygen (303), the 
union of sulfur dioxide and oxygen (617), and the oxidation of 
ammonia to nitric acid (570) . It finds extensive use catalytically 
in the manufacture of sulfuric and nitric acids. Since platinum 
has in recent years become much more valuable than gold, it has 
gained great popularity as a setting for jewels. Its use in jewelry 
was restricted and later prohibited during the war in order that 
the metal might be conserved for chemical uses. The use in 



536 Introduction to General Chemistry 

jewelry of this noble metal, so necessary for chemistry and in- 
dustry, ought permanently to be prohibited. The compounds of 
platinum are not of much practical importance. Until recently 
the chloride PtCl 4 , a very soluble yellow salt, was used in the 
quantitative determination of potassium. Potassium chloride 
forms with PtCl 4 a very difficultly soluble complex salt, potassium 
platinic chloride, K 2 PtCl6, into which all the potassium in a given 
sample could be converted and this form isolated and weighed. 
All platinum compounds are easily reduced to the free ele- 
ment. The chloride, for example, gives off chlorine upon 
being heated to a red heat, 

PtCl 4 ->Pt+2Cl 2 . 

819. Gold. — Gold (Au=i97.2) has been known from pre- 
historic times. It usually occurs in free form (native gold) but 
is also found in compounds such as the telluride (815). Several 
methods are in use for the extraction of gold from its ores. The 
simplest one is the amalgamation process, in which sand or 
crushed rock containing free gold is carried by running water 
over copper plates covered with a film of mercury. The gold 
forms an alloy (amalgam) with the mercury, from which it can 
be freed subsequently by distilling off the latter. In the chlori- 
nation process ores are extracted with a solution of chlorine, in 
which gold dissolves as auric chloride, AuCl 3 . From the 
chloride solution the gold may be precipitated by a variety 
of reducing agents, e.g., ferrous sulfate, etc. More important 
than the preceding is the cyanide process, by which enormous 
quantities of gold are recovered from low-grade ores. This 
process is based on the formation of readily soluble sodium 
auricyanide, NaAu(NC) 2 , which contains the very stable com- 
plex ion Au(NC) 2 ~. The reaction is as follows: 

4 Au+8NaNC-f-2H 2 0+0 2 ^4NaAu(NC) 2 +4NaOH. 

The crushed ore, which must sometimes be roasted to free it 
from sulfide, etc., is placed in immense tanks and extracted 
with a solution of sodium cyanide (665). The resulting extract 
is then run over zinc shavings to precipitate the gold, 

2 NaAu(NC) 2 +Zn->Na 2 Zn(NC) 4 +2Au . 






Some Additional Elements and Their Compounds 537 

Many copper, silver, and lead ores contain gold which 
accompanies the metals when they are obtained by smelting 
processes. This gold is recovered when the metals are refined 
electrolytically. 

The principal gold-producing regions in the order of their 
importance are the Transvaal (South Africa), United States 
including Alaska, Australia, Russia including Siberia, Mexico, 
and Canada. The value of the world's production of gold 
during the year 191 5 was 473 million dollars. 

Gold is a rather soft metal and is usually alloyed with copper 
or other metals. Pure gold is designated as 24 carat. Eighteen- 
carat gold, used in high-grade jewelry, is 75 per cent pure. 
United States gold coins consist of 90 per cent gold and 10 per 
cent copper. Gold is the most malleable and ductile of all 
metals. 

Metallic gold is not attacked by any of the common acids. 
It dissolves in aqua regia (562), a mixture of hydrochloric and 
nitric acids, to form auric chloride, AuCl 3 . This salt gives with 
hydrochloric acid, chlorauric acid, HAuCl 4 '3H 2 0, forming 
yellow crystals, and with sodium chloride, sodium chloraurate, 
NaAuCl 4 *2H 2 0. Aurous chloride, AuCl, is an insoluble salt 
obtained by gently heating AuCl 3 . 

A solution of auric chloride is used in toning photographs. 
In the ensuing reaction the silver of the picture is replaced (492) 
by metallic gold Au, 

AuCl 3 + 3 Ag->Au+ 3 AgCl . 

Two oxides of gold are known, aurous oxide, Au 2 0, and auric 
oxide, Au 2 3 ; these both give oxygen and metallic gold when 
heated. 

820. The Rare Earth Elements. — The term rare earth ele- 
ments is used to designate a large group of similar trivalent 
elements. The name indicates that these more or less rare 
elements bear some resemblance to aluminum, the silicate of 
which constitutes clay, a typical earthy substance. One of 
the most abundant sources of these elements is the mineral 
monazite, which is found in North Carolina, South Carolina. 



53 8 Introduction to General Chemistry 

Florida, Idaho, Brazil, and India. Monazite ordinarily occurs 
as a heavy yellow sand. It is a mixed phosphate of cerium, 
lanthanum, neodymium, and thorium, with smaller proportions 
of many other rare earth elements. 

Monazite sand is worked up primarily for the thorium it 
contains. Strictly speaking, thorium is not a rare earth element, 
although in its chemical behavior it is similar to these elements. 

Thorium nitrate, Th(N0 3 ) 4 , is extensively used # to make 
incandescent gas mantles; these consist, to the extent of over 
99 per cent, of thorium oxide, Th0 2 , and must contain o.6 to 
o.8 per cent of cerium oxide, Ce0 2 (301). 

A brief description of the chemistry of cerium, lanthanum, 
and neodymium will serve to illustrate the nature of the rare 
earth elements, which are all much alike in their properties. 
In free form the three elements mentioned are metals stable in 
air and toward water. An alloy of the three metals, called 
mixed metal, is made technically by the electrolysis of the molten 
chlorides CeCl 3 , LaCl 3 , NdCl 3 . An alloy of 70 per cent mixed 
metal and 30 per cent iron gives out a shower of sparks of burning 
metal when scratched with a file. It is for this reason known as 
pyrophoric alloy and finds use in cigar lighters, gas lighters, etc. 

These elements form bases and salts in which they are 
trivalent. The hydroxides, e.g., La(OH) 3 , are almost insoluble, 
moderately strong bases. The nitrates, chlorides, and sulfates 
are soluble and are not appreciably hydrolyzed in solution. The 
carbonates, phosphates, and oxalates are insoluble in water. 
Cerium differs from the other two elements in forming a series 
of 'quadrivalent salts derived from the oxide Ce0 2 . In these the 
eerie ion, Ce 4+ , is a strong oxidizing agent. The eerie- ion, 
Ce 4+ , is orange in color, while cerous ion, Ce +++ , is colorless, 
as is also lanthanum ion La +++ . Neodymium ion, Nd +++ , 
is a beautiful rose color. In addition to these rare earths 
there are eleven others, all of which are less common than the 
three here mentioned. 



CHAPTER XXXI 
CLASSIFICATION OF THE ELEMENTS. THE PERIODIC SYSTEM 

821. Introduction. — By reason of the enormous number of 
known chemical substances the best that a trained chemist can 
hope or even wish to do is to become thoroughly familiar with 
the principles and laws of his science and to know a moderate 
number of facts regarding the commoner substances, together 
with those with which his own field of work brings him in con- 
tact. His mind should be a laboratory and not a warehouse; 
otherwise he will soon find it so crowded with useless and unre- 
lated data (chemical junk!) that he has no ability to solve the 
new chemical problems that will constantly confront him. If one 
knows well the fundamental chemical principles and is familiar 
with the chemical behavior of the most typical of the elements 
he will find it easy to understand the chemistry of any unfamiliar 
element he may have occasion to study. Moreover, by reason 
of the relationships between the elements to be discussed in this 
chapter he will be able from a knowledge of a very few facts 
regarding the element in question to predict with more or less 
assurance the general behavior of the element and its com- 
pounds. 

We shall now consider briefly the entire list of the elements 
with the view of bringing out their relations to one another 
and of showing that they may be classified into groups (families) 
and series in such a way as clearly to exhibit these relationships. 
In fact, every element finds a definite place in this system of 
classification, which for reasons soon to be shown is called 
the Periodic System. 

822. Chemical Groups or Families. The Halogens. — The 
similarities in properties and behavior of chlorine, bromine, and 
iodine and of their corresponding compounds must be well 
known to the reader at this stage of advancement. Let us 
review some of the facts already established and consider also 

539 



54Q 



Introduction to General Chemistry 



a few additional ones. These elements are all non-metals which 
form the colorless gases HO, HBr, and HI, all of which dissolve 
easily in water to form strong acids. In these acids and their 
salts the halogen is always univalent. Although the free halo- 
gens are colored their ions in solution are all colorless. The salts, 
with the exception of those of silver, lead, and univalent mercury, 
are easily soluble in water. The three halogens being con- 
sidered all form oxygen acids and salts of a class represented by 
KCIO3. Thus we have the acids chloric, HC10 3 , bromic, 
HBr0 3 , and iodic, HI0 3 , and their corresponding salts, the 
chlorates, bromates, and iodates. 

If now we turn to the differences between the corresponding 
compounds of the three halogens we observe that the properties 
of bromine and its compounds are in most cases nearly inter- 
mediate between those of chlorine and iodine. Let us consider 
first the matter of atomic weights : 

Chlorine 35.5 

Iodine 127. 

Mean 81 .2 

Bromine 80. 

The atomic weight of bromine is seen to be very close to the 
mean of the atomic weights of the other two elements. The 
data shown in Table XLV illustrate the fact that bromine and 

TABLE XLV 

Comparison of the Properties of Chlorine, Bromine, and 
Iodine and Their Compounds 



Chlorine 


Bromine 


Gas 


Liquid 


Yellow 


Brown 


-102 


_o 

— / 


- 35° 


+ 59° 


i-5 


3- 2 


33 


75 


34 


67 


776° 


710 


79°° 


75o° 



Iodine 



Physical state 

Color of gas or vapor 

Melting-point 

Boiling-point 

Density. . .• 

Solubility of sodium salt* 

Solubility of potassium salt*. . . 
Melting-point of sodium salt. . . 
Melting-point of potassium salt. 



Solid 
Violet 
+ 114° 
+ 184° 
5-o 
180 
140 
650 
639° 



* Grams of salt dissolved by ioo g. of water at 20 



Classification of the Elements. The Periodic System 541 



its compounds stand nearly midway between chlorine and 
iodine and their corresponding compounds; so that if we know 
the properties of compounds of the latter we can predict pretty 
nearly those of the former. 

Fluorine, although classed as a halogen, is less closely 
related to the three elements just discussed than are the latter 
to one another. 

823. The Alkali Metals. — Sodium and potassium are called 
the alkali metals. This group or family also includes lithium 
(786), a much less common element, and in addition two rare 
elements, rubidium and caesium. These two elements form 
with sodium and potassium as closely related a group as do 
chlorine, bromine, and iodine. Lithium, although showing 
many points of similarity to the other four elements, bears about 
the same sort of relation to the others that fluorine does to the 
other halogens. Table XL VI gives some typical data on this 

TABLE XL VI 

Comparison of the Properties of the Alkali Metals and Their Compounds 



Atomic weights. 
Melting-points . 

Densities 

Atomic volumes 
Hydroxides. . . . 

Chlorides 

Nitrates 



Lithium 


Sodium 


Potassium 


Rubidium 


6.94 


23.00 


39.OI 


85-45 


186 


98° 


62° 


38° 


o.59 


O.97 


O.87 


1-52 


n. 9 


23-7 


44-8 


56.1 


LiOH 


NaOH 


KOH 


RbOH 


LiCl 


NaCl 


KC1 


RbCl 


LiN0 3 


NaN0 3 


KNO3 


RbN0 3 



Caesium 

132.81 
26° 

1.88 

70.6 

CsOH 

CsCl 

CsN0 3 



group. Again we find some simple relations between atomic 
weights, as in the case of the halogens. Thus the atomic weight 
of sodium, 23, is practically the mean of the atomic weights of 
lithium and potassium, while that of rubidium stands almost 
midway between those of potassium and caesium. Curiously 
enough, however, the atomic weight of potassium, 39, is much 
less than the mean of the values for the neighboring elements 
sodium and rubidium, namely 54. 

All of the five elements of this family are typical metals. 
As such they are all very active, caesium most so and lithium 
least. With water they all react readily to form strong, soluble 



54 2 Introduction to General Chemistry 

bases, resembling sodium hydroxide. This means that they all 
form positive ions, all of which are univalent. The metals are 
all soft (easily cut with a knife). They are all silver- white and 
have low melting-points, rubidium and caesium melting below 
blood heat. The melting-points decrease in order from lithium 
to caesium. The densities are all lower than those of any other 
metals. In general the higher the atomic weight the greater 
the density ; but we find that the value for sodium is exceptional 
in being a little higher than that of potassium. If, however, we 
calculate from the density and atomic weight the so-called atomic 
volume, that is the volume in cubic centimeters occupied by one 
gram atomic weight of the element, we find that these constants 
increase steadily with increasing atomic weight, and that the 
value for sodium is no longer irregular. Since the gram atomic 
weights (symbol weights) of all elements contain equal numbers 
of atoms, the atomic volumes are the volumes of equal numbers 
of atoms. By reference to Table XL VI we see that in this family 
the heavier an atom is the greater the volume it occupies. 

824. The Alkali Earth Group. — The alkali earth group 
includes the four elements calcium, strontium, barium, and 
radium. These are all bivalent metals whose hydroxides are 
moderately strong bases, and whose salts are typically repre- 
sented by calcium chloride, CaCl 2 , and sulfate, CaS0 4 . A 
detailed examination of this group would show that its members 
closely resemble one another in general, and also that a system- 
atic and gradual change of properties accompanies change of 
atomic weight just as in the halogen and alkali metal families. 

825. The Inert Gases. — The inert gases helium, neon, argon, 
krypton, and xenon (791-797) must obviously be considered as 
belonging to a separate family. They are all devoid of chemical 
properties and are therefore of valence zero in all cases. As 
gases their densities are of course all proportional to their 
molecular weights (Avogadro's Law, 193), The ratio of the 
molecular heat at constant pressure to that at constant volume 
is 1 . 66 in each case, thus showing that each gas is monatomic 
(793). The critical temperatures and boiling-points of the 
liquid gases all increase with increasing atomic weight from the 

/ 



Classification of the Elements. The Periodic System 543 

lowest values found in the case of helium. No other elements 
(excepting niton, radium emanation, chap, xxxii) resemble the 
five members of this family at all closely. Sometimes the inert 
gases are called the noble gases, because the most striking char- 
acteristic of the noble metals, gold and platinum, is the reluctance 
with which they enter into, chemical combination. 

826. Series of Elements. — If we arrange the elements in the 
order of their atomic weights the first of the series is obviously 
hydrogen. The next is helium. The next seven are shown in 
Table XL VII. The data presented in this table show that the 



TABLE XLVII 

The First Series of Elements 





Li 


Be 


B 


c 


N 





F 


Serial numbers 


3 

7 
1 86° 
LiCl 

1 


4 

9 
1800 
BeCl 2 

2 


5 
11 

2350° 

BCI3 

3 


6 

12 
3600 
CC1 4 

4 
CH 4 

4 


7 
14 

— 2IO° 


8 
16 

-218 


9 

19 

— 223 


Atomic weights 


Melting-points 


Chlorine compounds 

Valence toward chlorine . . 








Hydrogen compounds .... 


NH 3 

3 


H 2 

2 


HF 


Valence toward hydrogen . 








1 













valence toward chlorine increases by one as we pass from lithium, 
one, to carbon, four, and that the valence toward hydrogen 
decreases by units from carbon, four, to fluorine, one. There is 
also a gradual change of both physical and chemical properties 
as we go from lithium to fluorine. For example, lithium is a 
light, soft, silver- white metal (density 0.59). It readily decom- 
poses cold water. Lithium hydroxide is a strong base, and its 
salts are not hydrolyzed in solution. Beryllium is a white 
malleable metal of density 1 . 64. It does not tarnish in air, 
but it slowly decomposes hot water. Beryllium hydroxide is 
a rather weak base. The chloride is appreciably hydrolyzed in 
water solution. Boron (801-803) is a non-metallic solid of 
very high melting-point. It is not acted on by water or by 
oxygen except at high temperatures. Its hydroxide, B(OH),, 
or H3BO3, boric acid, is an extremely weak acid. The chloride 
BCI3 is completely hydrolyzed by water to form boric and 



544 



Introduction to General Chemistry 



hydrochloric acids. Carbon is strictly non-metallic in the form 
of diamond (the latter is an electrical, insulator-like glass), and 
only semimetallic as graphite (this has a semimetallic luster and 
conducts the current fairly well, 630). Carbon tetrachloride, 
CC1 4 (644), is a colorless liquid with none of the characteristics 
of a salt. The other three members of the series, nitrogen, 
oxygen, and fluorine, are all gases and all strictly non-metallic 
in character. Their activity increases very markedly from 
nitrogen, through oxygen, to fluorine. In their hydrogen com- 
pounds the last four elements of the series show a progressive 
decrease in valence from four for carbon to one for fluorine. 
Only the hydrogen compounds of nitrogen, oxygen, and fluorine 
form compounds with metals, e.g., NaNH 2 (527), NaOH, and 
NaF. The first is decomposed completely by water. 

827. The Second Series. — The seven elements beginning 
with sodium taken in increasing order of their atomic weights 
form a second series as illustrated in Table XLVIII. Again in 

TABLE XLVIII 

The Second Series of Elements 





Na 


Mg 


Al 


Si 


P 


S 


ci 


Serial numbers 


11 

23 

NaCl 

1 


12 

24-3 
MgCL 

2 


13 
27 
AICI3 
3 


14 

28 

SiCl 4 

4 
SiH 4 

4 


15 

3i 

PC1 S 

5 
PH 3 

3 


16 
32 


17 


Atomic weights 

Chlorine compounds 

Valence toward chlorine. . 


35-5 






Hvdrogen compounds .... 


H 2 S 

2 


HC1 


Valence toward hydrogen 






1 











this series metallic properties characterize the first three elements. 
The fourth, silicon, is a hard, black, difficultly fusible element 
with fair electrical conductivity; it may well be classed as semi- 
metallic. The following three elements are truly non-metallic. 
Chemical activity, great in the case of sodium, diminishes toward 
silicon, which is quite inert, and then increases again to a maxi- 
mum with chlorine. In this connection it is of importance to 
note that the activities of the first three metals are represented 
by their tendencies to take on positive charges (lose electrons, 
491), while those of the last three elements are due to a tendency 



Classification of the Elements. The Periodic System 545 

to acquire negative charges (gain electrons, 489). Furthermore 
the valences change by unity as we pass from one to the next 
element in the series. 

828. Correlation of the First and Second Series. — If we 
write the two series of elements in parallel columns a very 
remarkable relationship at once appears, as shown in Table 
XLIX. We see that Li and Na, both members of the alkali- 





TABLE 2 


CLIX 










First series 

Second series 

H compound 


Li 

Na 


Be 

Mg 


B 
Al 


C 

Si 

EH 4 

EC1 4 


N 
P 

EH 3 
EC1 S 




s 

EH 2 


F 
CI 
EH 


CI compound 


EC1 


EC1 2 


ECI3 











metal group, fall in the same vertical column; likewise the 
halogens, F and CI. Furthermore, if E is written as the general 
symbol of any element, the compounds with H or CI for the two 
elements of any vertical column can be represented by the same 
formula. In other words, this arrangement brings elements 
of similar properties into the same vertical columns. 

829. The. Periodic Table of the Elements. — Table L shows 
a systematic arrangement of 70 of the 83 elements. In this 
table the first two horizontal lines reproduce the two series of the 
preceding section and in addition include the elements helium, He, 
and neon, Ne. Following chlorine all known elements to and 
including cerium, Ce, 140.2, are given in the order of increasing 
atomic weights, with the exception that argon, A, 39 .9, precedes 
potassium, K, 39 . 1 ; and tellurium, Te, 127.5, precedes iodine, I, 
126.9. Between cerium, Ce, 140.2, and tantalum, Ta, 181 .5, a 
number of rare earth elements are omitted (820, 846). 

The first and second series, called also the first and second 
periods respectively, each includes eight elements. The 
double series of 18 elements beginning with argon, A, 39.9, 
and ending with bromine, 79.9, forms the third period. It 
will be noticed that a group of three elements, iron, Fe, cobalt, 
Co, and nickel, Ni, follows manganese, Mn, 54.9, in the column 
headed Group VIII. Copper, the next element after nickel, is 
placed in Group I and not under argon in Group 0. The 



546 



Introduction to General Chemistry 



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101.7 102.9 106.7 


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190.9 193. 1 195.2 




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P< xo 

CO 


00 
CO 


CM 
M 


A- 


o 

ft 

3 

s 

o 




£ O 


CO 
CO 


<N 

00 


CM 




I^H CM 
CM 


O 

4) 

Pi 


- 


CM 


CO 


* 


VO 


O 


^ 


saoinad 


ISOHS 




saora 


3d ON01 







Classification of the Elements. The Periodic System 547 

fourth period begins with krypton, Kr, 82.9, and ends with 
iodine, 126 .9. The fifth period includes the elements beginning 
with xenon, X, 130.2. The sixth period, which is defective, 
shows tantalum as its first and bismuth as its last member. 
The seventh period contains only four elements, of which niton, 
Nt, 222.4, is the first and uranium, U, 238.2, the last. No 
element of higher atomic weight than uranium is known. The 
first two periods are called short periods, the next four long 
periods, the seventh, and last, is fragmentary. 

830. The Nine Groups of Elements. The Zero Group. — The 
elements contained in any one of the nine vertical columns 
headed O to VIII constitute a group. There are, therefore, nine 
such groups. The zero group consists of the five inert gases 
(rare or noble gases studied earlier, 825), together with niton 
(radium emanation, see chap, xxxii), which is also inert chemi- 
cally. None of these gases forms any chemical compounds. The 
members of the zero group never have other than zero valence. 

831. The First Group. A and B Families. — Reference to the 
Periodic Table shows that five of the eight elements of Group I 
belong to the alkali-metals family. These elements are lithium, 
sodium, potassium, rubidium, and caesium. The other three 
elements, copper, silver, and gold, as we already know from 
their earlier study, are widely different from sodium and potas- 
sium, the best-known alkali metals. In many respects copper, 
silver, and gold are much alike and may be classed together as 
a family. Let us compare and contrast the two families of 
Group I, which we may designate as I A and IB respectively. 
The differences are set forth in Table LI. 

TABLE LI 





IA 


IB 


Elements 


Li.Na, K,Rb, Cs 


Cu, Ag, Au 


Density 


0.59 to 1 .88 
Very soft 
26° to 1 86° 
Give hydroxides 
Soluble strong bases 
Very high 


8.8 to 19.2 
Hard 


Hardness 


Melting-point 


960 to 1060 


Action of water 


No action 


Hydroxide 


Insoluble or unstable 


Displacement power (492) 


Very low 



548 Introduction to General Chemistry 

The differences are set forth in Table LI. The members of IA 
are uniformly univalent when in combination and form such 
compounds as EC1, EBr, EN0 3 , E 2 S0 4 , EOH, and E 2 0. In 
sodium peroxide, Na 2 2 , we believe that sodium is still univalent 
and represents the structural formula thus: Na • O • • Na (324). 
In IB copper and gold are variable in their valences; thus copper 
forms in addition to the commoner compounds in which it is 
bivalent (165) a series of univalent compounds, of which cuprous 
chloride, CuCl, is an example. Gold has so little chemical 
affinity that it does not form many stable compounds, but in 
these it is usually univalent or trivalent. It forms two oxides, 
Au 2 and Au 2 3 . If we hx our attention on the resemblances 
between copper, silver, and gold rather than on the divergences 
it is easy to see why they should be classed in the same family. 
The metals themselves are all permanent in air and not very 
active chemically; none of them sets free hydrogen from acids 
appreciably. They are moderately hard but are very malleable 
and ductile. They all melt in the neighborhood of 1000 . They 
are the best three conductors of heat and of electricity. They 
all form oxides, E 2 0, and their chlorides, EC1, are all white salts 
insoluble in water and dilute acids. It will be noted that Cu, 
Ag, and Au occur in the long periods 3, 4, and 6 respectively, 
and in each case in the second line of the period. Inspection 
will show the presence of A and B families in Groups II to VII 
but not in groups and VIII. 

832. The Second Group.: — Group II is made up of two 
families, IIA and IIB. The alkali earth elements Ca, Sr, Ba, 
and Ra are the typical members of IIA (824); while Zn, Cd 
(cadmium), and Hg constitute IIB. Beryllium, Be, and mag- 
nesium, Mg, in the first and second periods are somewhat 
more closely related in their properties to the family IIA 
than to the family IIB. The alkali earth family, IIA, is 
so called because of the resemblance of its members to 
the alkalies, IA, on the one hand, and to the earths, typified 
by aluminum in Group III, on the other. Metallic calcium 
is a moderately soft metal and is acted on rather rapidly by 
water to form a strongly basic hydroxide and hydrogen, but 



Classification of the Elements. The Periodic System 549 

by no means so rapidly as are the alkali metals. The hydroxides 
of IIA are not quite as strong bases as those of IA; but their 
salts with strong acids give neutral solutions. The three mem- 
bers of IIB, Zn, Cd, and Hg, are typical metals and are all unacted 
upon by water and are all untarnished by air. The hydroxides 
of IIB are weak, insoluble bases, and solutions of salts with strong 
acids are acid in reaction by reason of hydrolysis (436). Com- 
parison of IIA and IIB shows less difference in properties than 
was found in the case of IA and IB. With the exception of 
mercury all the elements of IIA and IIB form exclusively com- 
pounds in which they are bivalent. Typical formulae are EC1 2 , 
EBr 2 , E(N0 3 ) 2 , ES0 4 , EC0 3 , EO, and E(0H) 2 . Mercury is 
exceptional, since in mercurous compounds (333) it is univalent. 

833. The Third Group. — Of the nine elements composing the 
third group, only two, boron (801) and aluminum (174), are 
common. The others, with the exception of lanthanum, are 
rare. In their compounds the third-group elements are always 
trivalent and give products of the following types : E 2 3 , E (OH) 3 , 
ECI3, E(N0 3 ) 3 , E 2 (S0 4 ) 3 . 

Boron is the only member of the group that does not form 
salts with acids; its hydroxide, boric acid (802), has only weak 
acid properties. Aluminum hydroxide, Al(OH) 3 , is amphoteric 
(177), forming salts with both acids and bases; but these salts 
are all considerably hydrolyzed in solution, so that, for example, 
a solution of A1C1 3 reacts acid, while one of NaA10 2 reacts basic, 
to indicators (436) . 

Lanthanum hydroxide, La(OH) 3 , a white solid, is a rather 
strong base forming salts, e.g., La(N0 3 ) 3 , which are not hydro- 
lyzed in solution. - 

The rare earths (820), which doubtless should be considered 
as third-group elements, will be treated separately in a later 
section (846). 

834. The Fourth Group. — All the elements of the fourth 
group, with the exception of germanium (Ge=72.5), are of use 
technically. Carbon and silicon are of course the most impor- 
tant, and tin and lead come next. Titanum, zirconium, cerium, 
and thorium find interesting minor uses. The characteristic 



550 Introduction to General Chemistry 

valence of these elements, in their compounds, is four, so that 
the typical compounds have formulae like the following: E0 2 , 
E(OH) 4 , H 2 E0 3 , EC1 4 , E(N0 3 ) 4 , E(S0 4 ) 2 , etc. Only titanium, 
zirconium, tin, cerium, and thorium form salts of the last two 
types. Several of the members of this group form compounds 
in which they have a valence of two or three. Thus carbon and 
lead form the oxides CO (632) and PbO (167) respectively. In 
all stannous compounds (809) tin is bivalent, while in all its 
ordinary salts lead also has a valence of two. In the more stable 
salts of cerium, like cerous nitrate, Ce(N0 3 ) 3 , and cerous sulfate, 
Ce 2 (S0 4 ) 3 , this element is trivalent; but moist cerous hydroxide, 
Ce(OH) 3 , takes up oxygen from the air to form the more stable 
eerie hydroxide, Ce(OH) 4 . All the fourth-group elements form 
chlorides of the type EC1 4 . All of these except CeCl 4 and PbCl 4 
are stable toward heat; the chlorides of carbon, silicon, titanium, 
germanium, and tin are colorless liquids of low boiling-points 

(59° to 135*0- 

The hydroxides of the elements of the fourth group form only 
weak acids or weak bases. Carbon and silicon give carbonic acid 
and silicic acid (807) respectively; and titanium, zirconium, 
cerium, and thorium form hydroxides which are weak bases. 
Stannous and stannic hydroxides (809) and lead hydroxide, 
Pb(0H) 2 , are amphoteric. Titanium, zirconium, and thorium 
are the typical members of IVA; the position of cerium in this 
group is a bit uncertain. Germanium, tin, and lead compose 
IVB. Carbon and silicon bear resemblances to both of these 
families. Tin and lead are easily obtained as free metals, but 
it is doubtful whether even moderately pure metallic thorium 
has ever been obtained. 

835. The Fifth Group. — We have studied only five of the 
eight elements given in Group V, Table L. These are nitrogen, 
chapter xxi;. phosphorus, chapter xxiii; arsenic (810), antimony 
(811), and bismuth (812). The last three form family VB, to 
which the first two show points of similarity. Of the members 
of family VA only vanadium (V=5i) is of technical importance. 
It is a rather scarce element which finds an important use as a 
minor component of high-grade machine steel. Its common 



Classification of the Elements. The Periodic System 551 

oxide is V 2 5 , which is weakly basic and rather strongly acidic 
in character. The metavanadates, e.g., NH 4 V0 3 , correspond 
to the salts of metaphosphoric acid, HP0 3 (589). No technical 
uses have yet been found for the other two members of this 
family, the rare elements columbium (also called niobium) and 
tantalum. 

The following discussion will be confined to the members of 
VB, together with nitrogen and phosphorus. Of the five elements 
referred to, all but bismuth form gaseous hydrogen compounds 
of the type EH 3 . Of these only NH 3 forms with water a basic 
hydroxide. In these hydrides the fifth-group element is tri- 
valent. All five of the elements (including bismuth) form 
trimethyl derivatives, E(CH 3 ) 3 , corresponding to trimethyl 
amine (59). 

All five elements form trichlorides, EC1 3 , and P, As, and Sb 
form pentachlorides, EC1 S . All of the five elements being con- 
sidered form trioxides, E 2 3 , and all with the possible exception 
of bismuth form pentoxides, E 2 O s . Also, with the exception of 
bismuth, these elements form acids derived from their pentox- 
ides. These acids exist in various stages of hydration, of which 
HE0 3 is the most common. 

The highest valence of the elements of the fifth group is five, 
although in many cases the valence is only three. 

836. The Sixth Group.— The two families VIA and VIB are 
each made up of four elements. The first, VIA, consists of 
chromium (344), molybdenum (813), tungsten (814), and 
uranium; the second, VIB, embraces oxygen, sulfur (chap, xxiv), 
selenium (815), and tellurium (815). The members of VIA have 
a characteristic maximum valence of six, as exemplified by the 
oxides, E0 3 . These are all acidic in nature and lead to salts 
such as Na 2 E0 4 and Na 2 E 2 7 and even more complex formulae, 
e.g., Na6Mo 7 24 • 2 2H 2 0. In all these salts the sixth-group 
element is hexavalent. The members of family VIB all form 
hydrogen compounds, H 2 E, and these decrease in stability as the 
atomic weight of the sixth-group element increases. Water is 
extremely stable: hydrogen telluride decomposes with ease (,815). 
We should expect a normal valence of six for these elements; 



55 2 Introduction to General Chemistry 

but we find that for oxygen the ordinary valence is only two, 
although in some cases (not discussed in this book) a valence 
of four is very probable. Sulfur and tellurium form trioxides, 
E0 3 , and in these the two former are undoubtedly hexavalent? 
as they are also in the corresponding acids H 2 E0 4 and their salts- 
While selenium trioxide has not been made, the corresponding 
acid H 2 Se0 4 and its salts are well known (815). In them 
selenium has a valence of six. Sulfur, selenium, and tellurium 
also form oxides E0 2 and acids H 2 E0 3 and their salts. These, as 
we have learned, are converted by oxidation into the correspond- 
ing acids H 2 E0 4 or their salts. 

Contrasting VIA with VIB, we may say that while all the 
elements are acid-forming, only chromium and uranium are 
base-forming and give salts with acids. Chromium forms 
hydroxides in which it is trivalent. Thus Cr(OH) 3 is a weak 
base like Fe(OH) 3 . It also is capable of further reduction and 
gives a more strongly basic hydroxide, Cr(0H) 2 . The com- 
monest uranium salts are derived from the trioxide U0 3 . These 
salts may be considered as being formed from a compound 
U0 2 (0H) 2 , a diacid base about as strong as Fe(0H) 2 . Its 
reaction with nitric acid may be represented thus : 

U0 2 (OH) 2 +2HN0 3 ->U0 2 (N0 3 ) 2 + 2 H 2 . 

This salt is called uranyl nitrate. It forms large yellow crystals, 
U0 2 (N0 3 ) 2 • 6H 2 0. The corresponding uranyl chlorides, sul- 
fates, acetates, etc., are readily prepared. In these salts 
uranium is undoubtedly hexavalent. 

837. The Seventh Group. — Only one member of the A family 
of the seventh group is known; this is manganese (342). It is 
a hard metal the color of iron and does not tarnish in air. It 
does not form a hydride. In its common halogen compounds 
(true salts) it is bivalent. Its corresponding hydroxide is 
Mn(0H) 2 , an insoluble base which forms salts like MnBr 2 , 
Mn(N0 3 ) 2 , MnS0 4 , etc., and an oxide, MnO. Manganese also 
gives higher oxides, some of which are acid anhydrides. Potas- 
sium manganate, K 2 Mn0 4 , soluble green crystals, is a salt of the 
(unstable) manganic acid H 2 Mn0 4 , the anhydride of which would 



Classification of the Elements. The Periodic System 553 

be Mn0 3 . By oxidation (with chlorine, for example) solutions 
of nranganates give permanganates, 

2 K 2 Mn0 4 +Cl 2 -^2KCl+2KMn0 4 . 

Potassium permanganate (343) by treatment with sulfuric acid 
gives an explosive liquid (danger!) which is very probably man- 
ganese heptoxide, Mn 2 7 , the anhydride of permanganic acid, 
HMn0 4 . In the latter and its salts and anhydride manganese 
has a valence of seven. These compounds are all violent 
oxidizing agents. 

The halogen family, VIIB, is so well known and the close 
resemblances of its members have been so often referred to that 
further discussion of them would be superfluous. A few lines 
may be added touching the valence of chlorine, bromine, and 
iodine in their oxyacids and salts. Fluorine, it will be recalled, 
does not form such compounds. The other three halogens 
reach their maximum oxygen valence in the acids perchloric, 
HCIO4, perbromic, HBr0 4 , and periodic, HI0 4 . These may be 
considered as derived from hypothetical heptoxides, E 2 7 , in 
which the halogen has a positive valence of seven. There is 
thus an analogy between these elements and manganese in the 
peracids and their salts. 

838. The Eighth Group. — An inspection of the Periodic Table 
(829) will show that the arrangement of the members of the 
eighth group is different from that in any other group. In this 
group there are three lines of three elements each in place of the 
usual A and B columns. In first line we find iron, cobalt, and 
nickel, which are the only common elements of Group VIII. 
These elements have many similar properties. First of all it 
will be noted that their atomic weights are all close together. 
These three elements are all metals and all form salts of the types 
EC1 2 , E(N0 3 ) 2 , ES0 4 , E 3 (P0 4 ) 2 , and hydroxides and oxides 
E(0H) 2 and EO respectively. In all these compounds these 
three elements are bivalent. However, iron and to a lesser 
extent cobalt form hydroxides, oxides, and salts in which they 
are trivalent, e.g., E(OH) 3 , E 2 3) and EF 3 . The second line of 
the eighth group contains the three metals ruthenium, Ru, 



554 Introduction to General Chemistry 

rhodium, Rh, and palladium, Pd, all rare elements which closely 
resemble one another. Recently an alloy of palladium with gold 
has come into use as a substitute for platinum in the manufacture 
of crucibles and dishes for laboratory use. This alloy closely 
resembles platinum in its physical properties and in its inertness 
toward chemical reagents. The elements of the third line, 
osmium, Os, iridium, Ir, and platinum, Pt, form another sub- 
group in which the three members are much alike. All are 
extremely resistant to attack by most chemical reagents. 
Metallic osmium is of interest in that it has the greatest density 
of any known substance, namely 22.5. It is also extremely 
hard and melts only at the very high temperature of 2500 . 
It is also unique in forming an easily volatile and extremely 
poisonous oxide, Os0 4 , a solution of which, known as osmic acid, 
is a very important staining material for microscopic prepara- 
tions. In this oxide osmium has a valence of eight. 

839. Valence and the Structure of Inorganic Molecules. — 
The study of organic chemistry has made it very clear that the 
valence of each element of a compound can be definitely deter- 
mined only when the structure (648) of the molecule of the sub- 
stance is known. It is much easier to discover the structural 
formulae of organic compounds than of inorganic. There are 
two reasons for this: first, a single element, carbon, forms the 
backbone, so to speak, of all organic substances; and second, 
the enormous number of carbon compounds makes it possible 
to test theory by fact in a multitude of cases. The problem is 
very different with inorganic compounds, where we must deal 
with 80 or more elements, any one of which (excepting hydrogen, 
oxygen, and nitrogen) forms but few compounds in comparison 
with the host of carbon derivatives. Therefore when we attempt 
to show the structure of inorganic molecules the element of un- 
certainty is often great. 

In writing graphic formulae for the chlorides, oxides, hy- 
droxides, etc., of the elements of the first three groups we are 
on pretty safe ground, as these formulae are very simple. For 
common salt we have only one choice, Na« CI; similarly for other 
univalent elements. If for water (323) we write H-OH, for 



Classification of the Elements. The Periodic System 555 

sodium hydroxide we must write Na«0-H. Sodium oxide, 
Na 2 0, must be Na*ONa. In the second and third groups the 
task is nearly as easy. Calcium chloride is of course Cl-Ca-Cl, 
and the oxide must be simply Ca = 0, calcium and oxygen both 
being bivalent. Calcium hydroxide then is represented by 
H-OCa-OH. 

In the third group boron, which forms a trichloride, BC1 3 , is 
evidently trivalent. If boric acid, H 3 B0 3 , is written 

HCk 
HO— B 

ho/ 

the normal valences of all three elements are correctly repre- 
sented. Metaboric acid, HB0 2 , would then be HO*B=0 and 
sodium metaborate (802) NaO*B = 0. Sodium aluminate, 
NaA10 2 (177), would then be written NaO«Al = 0, which shows 
aluminum with its correct valence of three. 

The compound C(OH) 4 is not known, but carbonic acid, 
which would result from this by loss of water, is doubtless cor- 
rectly represented by 

HO v 

>C = 

HO/ 

which shows carbon with a valence of four. The structure of 

silicic acid is probably analogous. 

Nitrogen pentoxide, N 2 5 (555), is the anhydride of nitric 

acid. We might expect a hydroxide N(OH) 5 or H 5 N0 5 . By 

loss of water this could form first H 3 N0 4 and then HN0 3 . As 

a matter of fact a crystalline substance, HN0 3 -H 2 0, which is 

really H 3 N0 4 , is actually known (541). For this hydrate we may 

write the formula 

HOv 

HO— N = 

ho/ 



and for nitric acid itself 



HO-NC 

X) 



556 Introduction to General Chemistry 

Both formulae give nitrogen a valence of five. Orthophosphoric 
metaphosphoric acids probably have analogous formulae. 
For the two oxides of sulfur we write 

//° 
= S = OandO = Sf 

^0 

respectively, thereby assuming the valence of sulfur to be four 

in the first case and six in the second. The corresponding acids 

then become 

HOv H0 V /y 

>S = Oand )Sf 
HCK HCK X> 

The acids of the seventh group of the type HE0 4 very 

probably have the structure 

O 

II 
HO-E=0 

II 
O 

with a manganese or halogen atom of valence seven. 

Osmium, the only element of the eighth group which forms 
an oxide, E0 4 , doubtless has a valence here of eight. If so this 
oxide must be represented thus : 

O 

II 
0=0s=0 

II 

o 

840. Positive and Negative Valence. — We have already seen 
that the valence of an element is represented by the number of 
electrons each of its atoms loses or gains when it reacts to form 
a compound (484) . Thus when copper and chlorine unite (246) 
to form CuCl 2 each atom of copper loses two electrons and each 
atom of chlorine gains one. When sulfur burns to form S0 2 
(340) each atom of sulfur loses four electrons and each atom of 
oxygen gains two. But when sulfur and hydrogen combine to 
form H 2 S it is the hydrogen which loses electrons, since this 
element forms positive ions only (H + ), and therefore each atom 



Classification of the Elements. The Periodic System 557 

of sulfur must gain two electrons. The graphic formulae of the 
two sulfur compounds showing the distribution of charges 
resulting from the transfers of electrons are as follows: 



Sulfur dioxide, 
Hydrogen sulfide, 



— ++++-- 

o = s = o 

+ — + 
H— S— H 



In these formulae each plus sign indicates a loss of one electron. 
In S0 2 the valence of sulfur is four; in H 2 S it is two. But very 
plainly these are different kinds of valence, and we should dis- 
tinguish them by calling the first a positive valence and the 
second a negative valence. 

Other elements also exhibit both positive and negative 
valence. Thus phosphorus in PH 3 (588) has a negative valence 
of three and in PC1 3 (247, 576) a positive valence of three, while 
in PC1 5 it has a positive valence of five. 

TABLE LII 



Group 


IV 


V 


VI 


VII 


VIII 


Compound 


CH 4 

C0 2 , CCI4 

4 
8 


PH 3 

p 2 s ?pci s 

5 
8 


H 2 S 

S0 3 , H 2 S0 4 

6 
8 


HCl 
HC10 4 

8 




Negative valence 

Compound 



Os0 4 


Positive valence ....... 

Sum of +and— valence. 


8 
8 



A very remarkable law can be formulated from facts like 
those brought out by Table LII. For many elements of Groups 
IV to VIII the sum of the maximum positive and negative valences 
of an element is eight. 

841. Metals and Non-Metals. — We may now consider the 
way in which the metals and the non-metals are distributed in 
the Periodic Table (829). In the table the heavy line starting 
between beryllium and boron and ending with the boundary 
between the seventh and eighth groups separates all of the 
pronounced non-metallic elements from the others, which are 
all more or less metallic in character. The symbols of the former 
are printed in italics. It is true, however, that a few elements, 



558 Introduction to General Chemistry 

particularly chromium, molybdenum, and manganese, thus 
included with the non-metals, are in the free state definitely 
metallic and form salts with strong acids, e.g., Cr 2 (S0 4 ) 3 and 
Mn(N0 3 ) 2 . , But it is also a fact that these elements all form 
acids that yield with bases typical salts, e.g., Na 2 Cr0 4 , Na 2 Mo0 4 , 
and KMn0 4 , and in this respect the elements resemble the non- 
metals. 

It is particularly the non-metallic elements segregated by 
the heavy line of Table L that exhibit both positive and negative 
valence. The rare gases (Group O) have little in common with 
either the metallic or the non-metallic elements. 

842. Chemical Activity and the Periodic Table. — For prac- 
tical purposes we may define the chemical activity of an element 
as its tendency to enter into combination. If one element 
displaces another from the solution of one of its compounds 
(488-492) the first is the more active of the two. For example, 
in family IA the displacement order is Cu, Ag, Au, and this is 
therefore the order of the elements in respect to decreasing 
activity. For the halogen family, VIIB, the displacement order 
is F, CI, Br, I, and this again is the order of decreasing activity. 
A thorough study of the displacement tendencies and activities 
has established the following important generalization: In 
Groups I to VII, inclusive, in the A families the activity of an 
element increases with increasing atomic weight; in the B series 
activity decreases as the atomic weight increases. In accord with 
this law the greatest activities are found in the elements caesium, 
Cs, at the bottom of IA and fluorine at the top of VIIB. It is 
also important to note further that the valence of caesium in 
its compounds is positive, while that of fluorine is negative. 

On the other hand, the least active elements are met with at 
the bottom of Group VIII in the metals Os, Ir, and Pt, and at 
the top of Group O, where we find helium an element of no 
chemical activity and so little physical affinity (cohesion) that 
in liquid form it has the lowest boiling-point of all the elements 
and is the only element that has not been solidified. 

843. Periodic Properties. — We have now discussed, as far 
as the limits of this text will permit, the ways in which the 



Classification of the Elements. The Periodic System 559 

properties of the elements change as we pass from element to 
element in each family and from family to family and group to 
group in the table. It now remains to call attention to the 
remarkable manner in which the properties of the elements 
change as we pass through the table from period to period in the 
order of increasing atomic weights. The periodic fluctuation of 
properties is most easily exhibited by means of a diagram or 
graph. Let us take for illustration the variation of atomic 
volume (823) with atomic weight. In Fig. 116 the atomic vol- 
umes are plotted on the vertical axis and the atomic weights 
on the horizontal axis. The result is a very remarkable graph, 
in which each period is represented by a crude U-shaped 
curve. We see that there is a periodic repetition of like properties 
as we pass along the graph. If we should plot in a similar way 
any one of many other properties of the elements, for instance 
the melting-points or the compressibilities, we obtain periodic 
curves (graphs) of more or less similar forms. These facts are 
summarized by the statement first made by the great Russian 
chemist Mendelejeff: The properties of the elements are periodic 
functions* of their atomic weights. This means that the atomic 
weight of an element determines its properties. This is plain if 
we note that the place of any element in the Periodic Table is 
fixed by its atomic weight, and that the properties of the element 
are indicated by its position in the table. Among the properties 
of elements which show periodic relations we may mention 
valence, chemical activity, melting- and boiling-temperatures, 
conductivity for heat and for electricity, hardness, etc. 

844. History of the Periodic System. — The family relation- 
ships of the elements began to be discovered long before the 
Periodic Table as a whole was developed. As early as 1829 
Dobreiner pointed out the existence of triads of elements, such 
as CI, Br, I; Li, Na, K; and Ca, Sr, Ba, in which the atomic 
weight and other properties of the middle element of the triad 
were close to the mean of the other two. As time went on 
other families and their relations were recognized. The hrst 

1 The term function is much used in mathematics. If, for example, the value 
of x is dependent on the value of y, then x is said to be a function of v. 



560 



Introduction to General Chemistry 



attempt to arrange all of the elements in a single table was made 
by Newlands in 1864-66. A little later (1869-70) Mendelejeff 




published a table almost in the form of Table L. There were 
several places left vacant by Mendelejeff for elements as yet 



Classification of the Elements. The Periodic System 561 

undiscovered, and the whole zero group was of course absent, 
as none of these elements was discovered until many years later. 
It was Mendelejeff who pointed out many of the relations 
between the properties of the elements in essentially the form 
discussed in this chapter and drew the far-reaching conclusion 
that the properties of the elements are periodic functions of their 
atomic weights. This statement is known as the periodic law. 
The German chemist Lothar Meyer also discovered the periodic 
law independently and published a periodic table closely resem- 
bling that of Mendelejeff very shortly after the appearance of 
the latter's table. 

845. The Vacant Places of Mendelejeff's Table.— In Men- 
delejefPs table the spaces now occupied by scandium, gallium, 
and germanium were left vacant. He very boldly took the stand 
that these places represented elements still to be discovered. In 
fact, his description of these undiscovered elements, which he 
called ekaboron, ekaluminum, and ekasilicon respectively, was 
so complete and accurate that chemists knew just how they 
would behave if they were present in mixtures with other 
elements. In other words, they knew in just what sorts of 
chemical residues to expect these new elements, and it was not 
long until all three had been found. In 1875 Lecoq de Bois- 
baudran discovered ekaluminum and called it gallium (from 
Gaul, in honor of France); in 1879 Nilson discovered ekaboron 
and called it scandium (for Scandinavia); and in 1886 Winkler 

TABLE LIU 

Mendelejeff's Predictions, 187 i Winkler's Observations, 1886 

Ekasilicon ; Es = 7 2 Germanium ; Ge = 7 2 . 5 

Metal; density 5.5 Metal; density 5.47 

Oxide, Es0 2 ; density 4 . 7 Oxide, Ge0 2 ; density 4 . 7 

Chloride, EsCl 4 ; density 1.9 Chloride, GeCl 4 ; density 1.887 

Chloride should boil below ioo° Chloride boils at 86° 
Fluoride, EsF 4 ; not gaseous Fluoride, GeF 4 ; solid 

Sulfide, EsS 2 ; insoluble in water, Sulfide, GeS 2 ; insoluble in water, 
but soluble in ammonium sulfide but soluble in ammonium sulfide 
Tetra ethyl compound, Es(C 2 H 5 ) 4 ; Tetra ethyl compound, Ge(CaH s ) 4 ; 
density 0.96; boiling-point 160 density a little less than water; 

boiling-point 160 



562 Introduction to General Chemistry 

discovered ekasilicon, which he named germanium (for Germany) . 
In Table LIII some of the properties of the last-named element 
are set down opposite those predicted by MendelejefL 

MendelejerT also predicted the existence of ekamanganese 
with atomic weight about 100, but this element remains as yet 
unknown. 

When radium (480) was discovered by Mme Curie and its 
atomic weight was determined to be 226 it was at once placed at 
the bottom of IIA as a member of the alkaline earth family, of 
which its chemical properties (see chap, xxxii) show it to be a 
member. The zero family of the Periodic Table was added by 
Ramsay after the discovery of the inert gases (791-798) and 
served to round out the system in an unexpected but most 
satisfactory and suggestive fashion. 

846. The Rare Earths and the Periodic Table.— The fitting 
of the rare earths (820) into the Periodic Table presents a diffi- 
culty that is not yet very satisfactorily solved. These ele- 
ments, as we have seen, are all trivalent in their characteristic 
compounds and all much alike in their properties. They plainly 
form one large family quite as homogeneous as any other. Only 
the first member, lanthanum, La, and possibly also the second, 
cerium, Ce, fall into their proper places when these elements are 
arranged in the order of their increasing atomic weights. The 
problem of the position of the rare earths in the table is an 
interesting one that still awaits solution. 

847. The Position of Hydrogen in the Periodic Table. — As 
the Periodic Table is usually written hydrogen is not included. 
The chemistry of hydrogen is so unique that this element is in 
reality in a class by itself. Its positive valence of one suggests 
a place in the first group; but solid hydrogen (296) is not metallic 
and bears no physical resemblance to the elements of Group I. 
A critical discussion of the position of hydrogen in the periodic 
system would carry us beyond the scope of this text. 

848. Two Striking Anomalies. — In Table L there are two 
striking anomalies: argon with an atomic weight of 39.9 pre- 
cedes potassium with a value of 39.10; and tellurium, 127.5, 
precedes iodine, 126.9. Before the discovery of argon the case 



Classification of the Elements. The Periodic System 563 

of tellurium and iodine was unique, and for a long time chemists 
were inclined to think that the then accepted values of the atomic 
weights of these two elements were in error, for it seemed 
incredible that there should be an exception to a law that held 
so consistently for all other known elements. It was, of course, 
out of the question to place iodine before tellurium, for this 
would bring iodine in VIB with sulfur and selenium (815) and 
place tellurium among the halogens! It was possible that 
either or both of the elements used in early atomic-weight 
determinations were impure, or that methods of analysis were 
faulty, and that in consequence the atomic weight of iodine 
should be, in reality, greater than that of tellurium. The 
scientific importance of the problem thus presented made 
chemists realize the need of determining atomic weights of all 
elements with the greatest possible accuracy. As a result some 
of the world's most skilled chemists have given their best efforts 
to such work. In the cases of the two elements under discussion 
the final and universally accepted results are 

Tellurium = 127. 5 
Iodine =126.92 

and we are therefore forced to conclude that the position of an 
element in the Periodic Table is not of necessity strictly fixed by its 
atomic weight. The more recently discovered case of argon and 
potassium serves to confirm this conclusion. 

A discovery made by Mosely in 19 14 and discussed in the 
following chapter has thrown a new light on this whole matter. 
This discovery consists in nothing less than the finding of a 
property of the elements more fundamental than their atomic 
weights, namely, their atomic numbers. If the elements are 
arranged according to these numbers the two anomalies dis- 
appear and every other element falls in its correct position. 
But this is another story, which finds its logical place in the 
following chapter. 

849. The Newer Tables: The Harkins Table.— Since the 
time of MendelejefT many other tables have been arranged. 
We shall have space to present only one of these, namely, that 



564 Introduction to General Chemistry 

of Professor W. D. Harkins. This table should be of particular 
interest to those who expect to specialize in chemistry, since it 
shows the relationships between the elements much more clearly 
than does the Mendelejeff table. Fig. 117 is a projection of a 
space model of the table. 

The elements are arranged about a continuous helix in such 
a way that those of the same family are in the same vertical 
column. In the actual model the path of the helix is indicated 
by a heavy wire connecting balls which are lettered with the 
symbols of the elements which 'they represent. The families 
(831) are also connected by heavy vertical rods of brass, which 
thus form the supporting structure of the model. Each unit 
of atomic weight is represented by 1 cm. vertical distance on the 
spiral (see scale at the side of Fig. 117). The elements are 
arranged in the order of their atomic numbers. 

The helix begins with hydrogen, which is uniquely placed 
at the top of the table. Next the helix descends to the posi- 
tion of helium and makes a big loop, coming back to neon, which 
is of course just under helium. On this loop are found the 
elements of the first period of the Mendelejeff table, namely 
lithium, beryllium, boron, carbon, nitrogen, oxygen, and fluorine. 
After arriving at the position of neon the helix makes another 
big loop back to argon, and on this are the elements of the 
second period of the Mendelejeff table. These are sodium, 
magnesium, aluminum, silicon, phosphorus, sulfur, and chlorine. 
From argon the helix passes next to potassium, calcium, scan- 
dium, and titanium; then it begins the first inner loop, passing 
successively through the positions of vanadium, chromium, 
manganese, iron, cobalt, nickel, copper, zinc, and gallium before 
returning to the outer loop at the position of germanium. Next 
the helix takes a full outer turn through the positions of arsenic, 
selenium, bromine, krypton, etc. 

If we continue following the helix we find that the first four 
A-family elements and the last four B -family elements are on 
the outer loops, together with the zero group. »The second four 
A-family elements, the eighth group, and the first four B-family 
elements are on the inner loops. This brings corresponding A 



Classification of the Elements. The Periodic System 565 




Periodic Table by WD.Harkins 
Fig. 117 



566 Introduction to General Chemistry 

and B families together in a very ingenious way. At the bottom 
of the model a dotted line is traced to show the relative positions 
of the inner and outer loops. It will be noticed that the inner 
loops spring from the outer at the fourth group, and that from 
that point the distance between the two increases, becoming 
greatest between the zero and the eighth group. The rare 
earths are placed on the vertical column of IIIA, since their 
chemical properties show them to be closely related to this 
family. The last loop of the helix represents the positions 
of the radioactive elements, which will be discussed in the 
next chapter. 

This table has the advantage of representing the elements in 
a continuous series. It also shows the relationship of the A and 
B families more plainly than does the MendelejefT table. The 
vacancies in this table are for atomic numbers as yet unassigned 
to any element, and there is therefore a very strong probability 
that elements will be found to occupy these places. This was of 
course not true of many of the vacant spaces of the MendelejefT 
table. Professor Harkins counts hydrogen and helium as the 
first period and calls it the zero period. The subsequent periods 
he counts as beginning with the alkalies and ending with the 
noble gases. Thus, period one would begin with lithium and 
end with the neon. 

850. Summary. — In this chapter we have discussed a system 
of classification which brings all the elements into a single table 
(Table L, 829), in which the individuals are arranged in nine 
groups (vertical columns) and seven periods (horizontal rows). 
The first two periods are short ones, the others long (829) ; and 
in the latter the elements of Group I to VII are further classified 
into A and B subgroups or families (831) . In the first two (short) 
periods the elements of Groups I and II belong to the correspond- 
ing A families (831, 832), while in the same periods the members 
of Groups VI and VII are B-family elements (836, 837). The 
first- and second-period members of Groups III, IV, and V may 
be classed with either the respective A or B families. 

In the periodic system of arrangement most of the properties 
of elements and their compounds vary gradually as we pass 



Classification of the Elements. The Periodic System 567 

from one element to its neighbor in the table. This is true 
whether we pass from top to bottom of the vertical columns, 
considering each family separately, or pass through any one of 
the seven periods in the order of increasing atomic weights. 
Within a given family the properties change in a systematic 
manner in such a way that a knowledge of the chemical properties 
and the behavior of a given member can be deduced from those 
of each of its nearest relations. This is one of the most impor- 
tant facts about the periodic system. Furthermore in each 
period we find a recurrence, in like order but of different degree, 
of the properties that appear in all the periods. In consequence, 
if we plot the graph of any given property, taking the elements 
in serial order (in the order of increasing atomic numbers, 848), 
we get a periodic curve more or less closely resembling in form 
that for atomic volumes shown in Fig. 116. These facts are 
epitomized in the periodic law of MendelejefT: " The properties 
of the elements are periodic functions of their atomic weights." 

The simple and systematic way in which valence (both posi- 
tive and negative) changes from group to group (839, 840) is one 
of the most striking and useful facts brought out in this chapter, 
and one that should be firmly fixed in the mind of every chemist. 

Finally the Periodic Table is of great value to all students of 
chemistry in presenting to them a bird's-eye view of all the 
elements. In so doing it furnishes a mass of evidence for the 
unity of matter in the sense indicated earlier (470, 482), as will 
be discussed more fully in the following chapter. It seems to 
set a limit to the possible number of elements and also to indicate 
the nature of the yet undiscovered ones. Of more practical 
importance than all else the table makes possible a system of 
classification for an enormous number of chemical facts and thus 
permits through its mastery an orderly arrangement of the 
knowledge that constitutes the science of chemistry. 



CHAPTER XXXII 
RADIOACTIVITY AND THE NATURE OF MATTER 

851. Introduction. — The study of the history of chemistry 
shows very plainly that the rapid growth of chemical knowledge 
was in no small measure due to the innumerable attempts of 
the alchemists to convert the base metals into gold. Before the 
dawn of scientific chemistry belief in the possibility of trans- 
mutation of the elements was nearly universal. As time went 
on and the facts and ideas discussed in the early chapters of this 
book became known, alchemy merged gradually into chemistry, 
which soon became a science as well as an art. With the science 
of chemistry came the conviction that the elements were 
immutable kinds of matter, and that transmutation was a 
myth consisting of nothing more substantial than the hopes of 
the avaricious or the cupidity of charlatans. 

After some years of disrepute new life was injected into the 
old idea in 18 14 by a suggestion of Prout. This Englishman 
called attention to the fact that the atomic weights of the 
gaseous elements were whole numbers and therefore even mul- 
tiples of the atomic weight of hydrogen and expressed the opinion 
that the atoms of other elements were made up of atoms of 
hydrogen. This idea came to be known as Prout's hypothesis. 
The first defeat for the hypothesis came when it was proved that 
the atomic weight of chlorine was 35.5. Prout now claimed a 
half-atom of hydrogen as the unit. Later work on atomic 
weights (see Table XLI, 800) proved even this half -size unit too 
large if the hypotheses were to be applied to all elements. On 
the whole Prout had but few supporters. Nevertheless it is an 
incontestable fact that for many elements the atomic weights 
are much nearer whole numbers than can be accounted for by 
chance. 

The discovery of the periodic law furnished new and to 
many the most convincing evidence that all elements are but 

568 



Radioactivity and the Nature of Matter 569 

modifications of one primitive form of matter. But no satis- 
factory theory of the nature of matter and the interrelations of 
elements was developed until the knowledge of the main facts 
of radioactivity (480) had led to the disintegration hypothesis 
as their most plausible explanation. Since the discovery of 
radium and radioactive phenomena has furnished the key to 
one of nature's most profound mysteries we shall devote the 
present chapter to an account of this subject. The story of 
the discovery of radium is most instructive in illustrating the 
way in which scientific advances of the greatest significance often 
result from the thorough investigation of phenomena that seem 
trivial to persons who call themselves practical. We shall there- 
fore relate the facts in their historical sequence. 

852. The X-Rays and Phosphorescence. — The X-ray appa- 
ratus has already been described (476), and it has been men- 
tioned that the X-rays are produced when the cathode rays 
strike the target. It has also been stated that the cathode 
rays, non-luminous themselves, cause the glass walls of a cathode- 
ray tube to glow with a greenish-yellow phosphorescence. Thus 
there seemed to be a causal relation between phosphorescence 
and X-rays. This supposed relationship was made the subject 
of an extensive and thorough investigation by Henri Becquerel, 
professor of physics at the Sorbonne (University of Paris). It 
had long been known that a number of substances gave out a 
faint phosphorescent light for some time after exposure to any 
bright light. All such phosphorescent substances were carefully 
examined by Becquerel to learn if these glowing bodies gave out 
X-rays. The latter have the same effect on photographic plates 
that light has, but they can pass directly through lightproof 
paper. Accordingly Becquerel tested the effect of these phos- 
phorescent substances on plates wrapped in black paper. None 
showed any indication of the production of X-rays, with one 
exception: this was potassium uranyl sulfate, K 2 UO_,(S0 4 ) 2 • 4ILC), 
a salt of the rather uncommon element uranium (836). 

853. The Becquerel Rays. — Not only did the above- 
mentioned uranium salt give out rays that, like X-rays, pene- 
trated black, lightproof paper and acted on a photographic 



570 Introduction to General Chemistry 

plate while phosphorescing after it had been exposed to sun- 
light, but the salt showed equal photoactivity if prepared and also 
tested wholly in the dark, under which condition it did not phos- 
phoresce. Furthermore all other uranium compounds tested 
showed similar photoactivity, although they were not phos- 
phorescent even if previously exposed to light. Obviously 
these new rays, called at first Becquerel rays, were not caused 
by phosphorescence but were produced by the element uranium 
in all forms of chemical combination. 

854. Ionization of Air by Becquerel Rays. — It was also soon 
discovered that Becquerel rays ionize air (770), so that by 
means of the discharge of a gold-leaf electroscope the activity 
of a substance could be tested much more readily than by 
the photographic method. A convenient form of electroscope 
suitable for projection-lantern experiments is shown in Fig. 109 
(770). Two parallel sides are made of glass plates; the rest of 
the case is of metal. A wire passes through the amber plug 
insulator and ends in a brass strip 6 to 8 mm. wide and 5 or 6 cm. 
long. A single leaf of gold or aluminum is attached to the brass 
strip, as shown in the figure. The electroscope may be charged 
by means of an ebonite rod, celluloid comb, fountain-pen holder, 
etc., electrified by being rubbed with a piece of woolen cloth. 
The uranium compound is held within 2 or 3 cm. of the end of 
the wire carrying the gold-leaf system. The rate of discharge 
of the electroscope under fixed conditions is a measure of the 
intensity of the Becquerel rays. 

855. The Discoveries of Mme Curie. — Becquerel's dis- 
covery 'of the rays known by his name was published in 1896, 
about two years after Roentgen had discovered X-rays. At 
about this time Becquerel suggested to Mme Curie, the brilliant 
wife of Becquerel's colleague, Professor Curie, that the investiga- 
tion of natural uranium minerals should prove interesting. 
Mme Curie's experiments soon showed that all uranium minerals 
were photographically active; they were considerably more 
active in ionizing air than one would expect from their uranium 
content. This fact seemed to suggest that the larger part of 
the activity of the minerals was eliminated during the chemical 



Radioactivity and the Nature of Matter 571 

treatment used in making the pure uranium compounds. Search 
was therefore made for other active substances among the 
by-products of the preparation of uranium salts; whereupon it 
was discovered that the small amount of barium extracted from 
the waste materials was many times as active as any pure 
uranium salt, notwithstanding the fact that ordinary barium 
salts were entirely inactive. 

856. The Discovery of Radium. — After considerable quanti- 
ties of active barium chloride had been extracted from uranium 
mineral residues Mme Curie undertook to separate from the 
barium the new substance to which she attributed the activity. 
She found that if a hot, saturated solution of the active barium 
chloride was allowed to cool about half of the dissolved salt 
crystallized out, but these crystals contained 70 per cent or 
more of the active material. Thus the crystals were appre- 
ciably richer in the active matter than the original material. 
By repeated recrystallization of the crystals that separated in this 
way Mme Curie finally obtained a product that was thousands 
of times more active than an equal weight of a uranium salt. 
This product was still largely barium chloride; but that it also 
contained a new element was shown conclusively by the fact 
that its spectrum (786) contained in addition to those of barium 
several lines not belonging to any known element. It was 
therefore certain that a new element had been discovered, and 
it seemed probable that it was this element which had produced 
the powerful Becquerel rays that caused the photoactivity and 
ionizing power of the material. On account of its ability to 
send out rays or radiations Mme Curie named the supposed 
new element radium. 

The next problem was to free the radium salt from the accom- 
panying barium. This was finally accomplished by a long series 
of crystallizations of the chlorides of the barium-radium mixture. 
Finally pure radium chloride free from barium was obtained. It 
was a white crystalline salt closely resembling barium chloride 
in appearance. Its activity, called now radioactivity, was two 
or three million times as great as that of an equal weight of a 
uranium salt. 



572 Introduction to General Chemistry 

857. The Atomic Weight of Radium and Its Position in the 
Periodic Table. — If an element does not form volatile com- 
pounds its atomic weight cannot be determined by means of the 
law of minimum weights (64) ; and if it is not obtainable in con- 
siderable quantity in uncombined solid form the method based 
on the law of Dulong and Petit (229) is also unavailable. One 
might make use of a method based on direct or indirect deter- 
mination of the osmotic pressure of a solution (717-722), but 
this involves a complication in case an ionized salt is used. As 
none of the ways just mentioned could be applied in the case of 
radium, of which but a fraction of a gram was available and that 
worth thousands of dollars in its cost of preparation, other means 
of fixing its atomic weight were used. 

If we analyze a chloride of a new element and find its per- 
centage composition, one additional fact must be known in 
order to fix the atomic weight of the element, namely, its valence. 
But if we can discover to what group of the Periodic Table the 
new element belongs we can of course infer its valence (839). 
We have seen that radium resembles barium so closely, in some 
respects at least, that the separation of the chlorides of the two 
elements is very difficult. This indicates that radium is a mem- 
ber of the alkali earth group containing calcium, strontium, and 
barium (824, 832). Furthermore it was found that radium also 
resembles barium in the following respects: its carbonate is 
insoluble in water, its sulfate is insoluble in water and in strong 
acids, and its hydroxide is soluble, since radium chloride solution 
gives no precipitate with either ammonia or sodium hydroxide. 
Furthermore the free element cannot be deposited on a platinum 
electrode from a water solution by electrolysis. It therefore 
seemed very probable that radium was a second-group element 
and had a valence of two. The analysis of the pure chloride 
was made by Mme Curie ; the result, with an assumed valence of 
two, led to a calculated atomic weight of 226. Since the space 
in the Periodic Table (829) corresponding to this atomic weight 
was vacant, no doubt remained of the correctness of the atomic 
weight assigned to the new element. 



Radioactivity and the Nature of Matter 573 

858. The Alpha Rays. — It has already been stated that 
radium emits three kinds of rays, designated as alpha, beta, and 
gamma, although, strictly speaking, only the last named is a true 
radiation (that is, a vibration in the luminous ether), since the 
alpha rays are atoms of helium (795) and the beta rays electrons 
(480). Entirely reliable experiments, which are too intricate 
to be described here, have shown that the alpha particles, as we 
shall now call them, are shot out from radioactive atoms with 
velocities of 5,000 to 12,000 miles a second. They travel in 
straight lines through air for distances (called their ranges) of 
3 to 8 cm. before they slow down to the average velocity of air 
molecules (about one-fourth of a mile a second, 197). The 
alpha particle of radium has a range of 3.5 cm., and in this 
distance it collides with many thousands of air molecules; but 
instead of rebounding as an elastic body and changing its direc- 
tion after each collision the helium atom constituting an alpha 
particle passes straight through the air molecule, as a rule, with- 
out changing its direction of motion. This fact furnishes the 
most convincing evidence that molecules are not solid but are 
made up of a group of widely spaced particles, the electrons (470, 
483). It is the alpha particles which cause the major part of 
the ionization of air exposed to radium rays. Each alpha particle 
leaves in its wake thousands of ions : each ion is an air molecule 
from which an electron has been detached, leaving a positive ion ; 
or to which a detached electron has united, forming a negative 
ion. All air ions carry single electric charges only. Alpha 
particles are easily stopped by thin sheets of any solid substance, 
such as paper, glass, or metal. 

Proof that the alpha rays are helium atoms was conclusively 
furnished in the following way: Radium, contained in a sealed 
glass tube with walls so thin that the alpha particles could barely 
pass through, was surrounded by a sheet of lead which completely 
stopped the particles. After several hours the lead was removed 
and sealed in a larger glass tube having wires for the passage of 
electric sparks. The lead was then melted by heating the glass 
tube; this set free the helium atoms that were imbedded in the 



574 



Introduction to General Chemistry 



lead, as was proved by the fact that on passing an electric dis- 
charge through the tube the spectrum of helium was easily 
recognized by observation with a spectroscope (785). 

859. Luminous Effect of the Alpha Rays. — When alpha rays 
strike a screen covered with a layer of a specially prepared form 
of crystalline zinc sulfide, ZnS, a pale phosphorescent light is 
produced. Luminous watch dials, which have recently become 
so popular, are produced on this principle. From one to ten 
parts of radium to ten thousand parts of zinc sulfide are used in 
making the luminous paint for such purposes. Radium luminous 




Fig. ni 



paint was extensively used in the war for making luminous dials 
for aeroplane instruments, submarine instruments, trench com- 
passes, gun sights, etc. 

If one views in the dark, by means of a good lens, a zinc 
sulfide screen exposed to alpha rays from a minute amount of 
radium a very beautiful effect will be seen. Instead of a uniform 
glow sparks of light will be seen appearing here and there 
at random over the screen. It is now certain that each tiny 
flash of light is the result of the striking of a single helium atom. 
The simple device consisting of a zinc sulfide screen, a bit of 
radium, and a lens was called a spinthariscope by Crookes (477), 
its inventor. The same effect may be seen if the radium paint 



Radioactivity and the Nature of Matter 



575 



on a luminous watch dial is examined in the dark with a powerful 
lens. 

860. The Beta Rays. — The beta rays are electrons (480) shot 
out from radioactive atoms with velocities closely approaching 
the velocity of light. The beta rays are much more penetrating 
than the alpha rays and can pass through ten to twenty sheets 
of paper. It is chiefly these rays (and in small measure the 
gamma rays) that are photographically active. Fig. 118 shows 
a photograph taken by radium rays. The photographic plate 
was wrapped in lightproof paper; the objects shown in the 
figure were laid on the paper and the whole exposed to the 
radium at a distance of 10 cm. for a few minutes. 
The plate was then developed in the usual way, and 
from the negative so obtained the positive print 
reproduced in the figure was made. 

A curious radium clock has been constructed by 
R. J. Strutt, an English physicist, to show the con- 
tinuous loss of negative electricity as electrons by 
radium. The radium is inclosed in a small tube, A, 
Fig. 119, made of glass thick enough to prevent the 
alpha rays from escaping but not thick enough to 
stop the beta rays. This small tube is attached at 
the top to an insulating rod, B, of quartz and carries 
at its lower end a pair of gold leaves, C. The tube 
and gold leaves are fixed in a glass vessel having a 
pair of tin-foil strips, D, pasted on the inside opposite the gold 
leaves.. These strips are electrically connected to a wire sealed 
through the glass and intended to furnish a connection to the 
earth. The larger glass vessel is highly evacuated. The action 
of the apparatus is as follows : The beta rays (electrons) from the 
radium tube escape to earth by way of the tin-foil strips, D. 
This leaves the radium tube with a deficiency of negative elec- 
tricity and therefore positively charged. The positive charge 
causes the gold leaves to diverge (by reason of the repulsion of 
like charges) until they touch the tin-foil strips, when they are 
at once discharged and fall together. Then the whole cycle 
begins again. Since the discharge of the gold leaves occurs at 




119 



576 Introduction to General Chemistry 

regular intervals Strutt called the device a radium clock. If the 
vessel were not evacuated ionization of the air would occur and 
conduct electricity away so fast that the gold leaves would not 
become charged. The beta rays can also ionize gases but are 
less effective than the alpha rays. 

861. The Gamma Rays and X-Rays. — The gamma rays 
closely resemble X-rays. They are far more penetrating than 
the beta rays and can pass through several millimeters of metal 
with little absorption and are only completely absorbed by several 
centimeters of lead, which is more effective in stopping these 
rays than any other common substance. The gamma rays 
ionize gases and also affect a photographic plate. Several 
minerals glow perceptibly when exposed to the gamma rays of 
a considerable quantity of radium (say ioomg.). Willemite, a 
silicate of zinc (807), is especially active in this way. Diamonds 
also glow with a clear white light under the gamma rays, so that 
these rays may be used practically to distinguish a real from an 
imitation diamond. 

The true physical nature of gamma rays and X-rays was for 
a long time in doubt. The critical work which cleared up the 
question will be discussed later (877). 

862. Heat from Radium. — Not long after radium had been 
obtained in considerable quantity it was discovered that a tube 
of this unique element was somewhat warmer than its surround- 
ings. Subsequent investigation by Rutherford (481) showed 
that one gram of radium gives out 134 calories per hour and 
keeps up this rate of heat production continuously with appar- 
ently little change. 

863. Radium Emanation. — The ionization of a gas does not 
occur spontaneously, as does the ionization of an electrolyte. 
Moreover, very quickly after the ionizing agent (say a tube of 
radium) is removed the ions of a gas neutralize one another's 
charges and the gas becomes nonconducting. With this fact 
in mind let us consider the following curious observation: If a 
current of air is led over a radium compound, or better through 
a solution of a radium salt, the air will become strongly ionized 
and remain so for hours and even days, losing its electrical con- 



Radioactivity and the Nature of Matter 577 

ductivity so slowly that only after about four days is half of 
the effect lost. In fact, it seemed as if the air had taken up 
radioactive matter from the radium, since this active air also 
caused a zinc sulfide screen to glow (859), and this glow was 
found by examination with a lens to be the result of scintillating 
flashes like those caused by alpha rays. In consequence the 
radium was said to have given off a radioactive emanation to the 
air. The material nature of the emanation was first clearly 
shown by experiments by Rutherford, in which air containing 
emanation was passed through a metal tube cooled externally 
by liquid air (776). Every trace of emanation was removed 
from the air that passed through the cooled tube; but when' the 
latter was removed from the liquid air and allowed to come 
to room temperature the whole of the emanation was recovered 
by blowing fresh air through the tube. The explanation of 
this experiment given by Rutherford is very simple: Radium 
emanation is a gaseous radioactive substance which condenses 
practically completely at the temperature of liquid air. Upon 
becoming warm it again volatilizes. This explanation has 
proved to be correct in every particular. 

864. The Formation and Decay of Radium Emanation. — The 
whole of the emanation contained in a solution of a radium salt 
can be removed by blowing a current of air through the solution 
for a few minutes. If the solution is sealed air-tight, new emana- 
tion will gradually accumulate in it. After about 40 days there 
is present in the sealed vessel a maximum of emanation strictly 
proportional to the quantity of radium present. The emanation 
accumulates at a regular rate, as follows: after 3.85 days 
from the time the solution was sealed (after all emanation has 
previously been removed) half of the maximum quantity of 
emanation is present; after 2X3.85 days, three-fourths of the 
maximum; after 3X3.85 days, seven-eighths of the maximum, 
etc., so that after 40 days over 99 .9 per cent of the maximum is 
present. We have seen that half of the activity of a given 
portion of emanation is lost in about 4 days (863). Strictly 
speaking, exactly half is lost in 3 .85 days; in 2X^ .85 days half 
of the remainder is lost, or a total of three-fourths; in 3X3 .85 



578 



Introduction to General Chemistry 



days half of the remainder is again lost, or a total of seven- 
eighths, etc. The emanation is said to decay to half value in 
j . 83 days. The interval of 3 . 85 days is called the period of half 
decay, or briefly the period. The rates of formation and decay 
of radium emanation are shown by the two graphs of Fig. 120. 
The proportion of emanation contained in a sealed sample of 
radium after an interval of 40 days is a maximum, because from 

this time on the decay of the 
emanation present exactly 
compensates the formation 
of new emanation. 

865. Theory of Radio- 
active Change. — The many 
facts already cited show 
that radium has properties 
that place it in a class by 
itself. For a number of 
years after the discovery of 
radium the cause of its 
conspicuous of all scientific 
scientists declared that the 




TIME IN DAYS. 



Fig. 



strange behavior was the most 
mysteries. Some faint-hearted 
known facts disproved the two most fundamental scientific 
laws: the law of the conservation of matter and the law of the 
conservation of energy. It was also said that a serious incon- 
sistency existed in the statement that radium is an element, if 
it is really true that radium gives rise to helium and radium 
emanation. The epoch-making hypothesis which fairly met 
and explained every known fact and withstood the most 
searching criticism was one of the boldest and at the same time 
the simplest and most comprehensive ever introduced into our 
science. This was the disintegration hypothests of Rutherford 
and Soddy (481). These scientists were at the time professors 
in the department of physics of McGill University, Montreal. 

The sketch of this remarkable hypothesis already given (481) 
should now be read again. According to this hypothesis radio- 
active atoms are more or less unstable systems. The first step 
in the disintegration of a radium atom is the throwing off of an 



Radioactivity and the Nature of Matter 579 

atom of helium (an alpha particle) and the leaving behind 
of a residual atom of smaller mass, an atom of radium emana- 
tion (Em), 

Ra->He+Em . 

If we suppose that the helium atom shot out was in rapid orbital 
motion while it formed a part of the radium atom we have at 
once a simple explanation of the high velocity of the alpha 
particle and therefore of the energy given out by radium as 
heat (862), since the kinetic energy of the alpha rays must be 
changed into heat when the rays are stopped (368-371). 

866. The Radium Series. — A solid body of any kind exposed 
to radium emanation for some time becomes itself strongly radio- 
active, giving out alpha, beta, and gamma rays. This so-called 
excited activity decays rather rapidly (to half in about thirty 
minutes) . Elaborate experiments have shown that the explana- 
tion of these facts is as follows: Just as a radium atom disin- 
tegrates, throwing off an alpha ray (helium atom) and leaving 
an atom of emanation, so an atom of the latter in turn shoots out 
another helium atom and leaves a new and lighter residual atom 
of a substance called radium A, having a very short period (864) . 
Radium A also gives out alpha rays and forms radium B, and the 
latter in turn passes into radium C. The series of radioactive 
products of radium is therefore as follows: 

Ra^Em^RaA^RaB->RaC . 

The excited activity obtained from the emanation consists 
of the products A, B, and C. These, together with the emana- 
tion, are present in radium sealed so as to retain the emanation. 
Radium itself gives only alpha rays. The beta and gamma rays 
come largely from radium C. 

867. The Origin of Radium. — Soon after the discovery of 
the nature of radioactive change attempts were made to discover 
the rate of decay of radium itself. Several methods were 
devised. The simplest consisted in counting the number of 
alpha particles given off by a known amount of radium in a 
measured period of time. It was arranged so that each alpha 
particle would strike a zinc sulfide screen and produce a flash of 



580 Introduction to General Chemistry 

light (859), the number of which flashes could be counted. 
Since we know the atomic weight of radium (857) and the actual 
number of atoms in one gram atomic weight of any element 
(6.06X10 23 ) (194), such an experiment furnished a means of 
calculating the rate of decay of radium. This rate is about 
0.04 per cent per year, which corresponds to a period of 1,850 
years. But even this period, long as it is in terms of the span 
of human life, is very short compared with the age of the earth 
as judged by the conclusions of geologists; so that if the earth 
were at the beginning made of pure radium, after even a million 
years far less radium would remain than we now find in many 
radium-bearing minerals. But as the age of the earth is to be 
reckoned in hundreds of millions of years, and as there is a 
practical certainty that the proportion of radium never was 
high, we are forced to the conclusion that the radium now found 
in the earth is being formed as fast as it decays. When these con- 
clusions had been reached it became a matter of much scientific 
interest to discover the origin of radium. 

The parent of radium, as we may call the hypothetical sub- 
stance from which we may suppose radium to be formed, would 
very likely conform to the following specifications: It would be 
associated with radium in minerals; the ratio of radium to the 
substance would be constant (for the same sort of reason that 
sealed radium contains, after a sufficient interval, a constant 
proportion of emanation); the parent would probably be radio- 
active, with a very slow rate of decay compared with that of 
radium; the atomic weight of the parent would exceed that of 
radium. Only one element conforms to all these specifications, 
namely uranium (836). This element is beyond doubt the parent 
of radium. The latter element is found only in uranium minerals 
and in the fixed ratio of one part of radium to three million parts 
by weight of uranium. 

Uranium-bearing minerals always contain mechanically 
imprisoned helium which has been formed by radioactive 
change. In fact it was in such a mineral, cleveite, that helium 
was discovered (794). 



Radioactivity and the Nature of Matter 



581 



868. The Uranium Series. — The change of uranium into 
radium is not direct but takes place by the intermediate forma- 
tion of a series of radioactive products, as indicated in Table LIV, 
which also includes the entire radium series. 



TABLE LIV 

The Uranium-Radium Series (Slightly Condensed) 



Element 


Period 


Rays 


Atomic Weight 


Uranium 

Uranium X x 

Uranium X 2 

Uranium 2 

Ionium 


5X10 9 years 
24 . 6 days 
69 sec. 
2X10 6 years 
2X105 years 
1,850 years 
3 . 85 days 
3.0 min. 
26.8 min. 
19.5 min. 
12 years 
5 days 
136 days 
? 


a 

a 
a 
a 
a 
a 

a, jS, T 


238 
234 
234 
234 
230 
226 


Radium 


Emanation. 

Radium A 

" B 

" C 

D 


222 
218 
214 
214 
2IO 


" E 

" F 

G 



a 


2IO 
2IO 
206 









869. The Atomic Weights of Radioactive Substances. — If 

the disintegration hypothesis is correct the atomic weight of a 
product formed by an alpha-ray change should be less than that 
of its parent by the atomic weight of helium, namely 4 units. 
By reason of the very small mass of an electron (479) the atomic 
weight is practically unchanged as the result of a beta-ray 
change. Taking the atomic weight of uranium (836) in round 
numbers as 238, the calculated atomic weights of its products 
are those given in the last column of Table LIV (868). We can 
easily test these conclusions in one case, that of radium. The 
calculated value is 226.2, if U= 232.2, while Mme Curie found by 
experiment the nearly identical value 226. Therefore it appears 
probable that the calculated values given in the table are 
essentially correct. 

870. Thorium and Its Active Products. — The element thorium 
(820, 834) is radioactive with an intensity of activity about equal 
to that of uranium. Like the latter, the former gives rise to a 



582 Introduction to General Chemistry 

long series of active products, of which the first, mesothorium, 
is the only one of technical importance. Mesothorium is identical 
with radium in its chemical behavior, and no means is known of 
separating a mixture of the two. They are readily distinguished 
by their periods and by their radioactive products. Meso- 
thorium has a period of 5.5 years as against 1,850 years for 
radium. Mesothorium itself gives no rays but changes into a 
series of active products giving alpha, beta, and gamma rays. 
Each product has its own characteristic chemical properties, 
rays, and period. One of these products is radiothorium, having 
a period of two years. This gives out alpha rays and forms 
thorium X, which in turn forms a gaseous emanation having a 
period of less than one minute. Mesothorium is produced as 
a by-product of the manufacture of thorium and is used as a 
substitute for radium. 

871. Isotopes. — A most unlooked-for set of facts was brought 
to light by the investigation of the chemical behavior of certain 
radioactive products. We have mentioned that mesothorium 
and radium are identical in chemical behavior (870). Thorium 
X is also identical chemically with these two, and no means is 
known of separating mixtures of this substance and either 
radium or mesothorium. Substances identical in chemical 
behavior are called isotopes. 

Several groups of isotopes are known. Thus radiothorium 
and uranium X are isotopic with thorium, since in every chemical 
reaction they both behave qualitatively and quantitatively 
exactly like thorium. Each of the three can be obtained 
separately, but if we mix the separate substances it is impossible, 
by any known means, to separate them. The final product of 
the radium series, formed by the decay of RaF and known as 
RaG (at. wt. 206), is an inactive substance isotopic with lead 
(at. wt. 207.2), and curiously enough RaD (at. wt. 210) is also 
isotopic with lead and therefore also with RaG. 

Isotopes are not identical in all their properties; otherwise 
they would be identical substances and consequently indis- 
tinguishable. Thus, for example, radium differs from meso- 
thorium in the following respects: Radium has a period of 



Radioactivity and the Nature of Matter 



583 



1,850 years; mesothorium has a period of 5.5 years. Ra gives 
alpha rays; Ms gives no rays. Ra changes into a gas, Em, 
which gives alpha rays and has a period of 3.85 days; Ms 
changes into a solid which gives beta rays and has a period of 
6.2 hours. Ra has an atomic weight of 226.0, while Ms has 
the value 228.4. No difference in the spectra of isotopes has 
yet been discovered. 

872. The Valence of Radioactive Substances. — Although but 
very few of the radioactive products of the uranium-radium 
series (Table LIV) or the thorium series have as yet been obtained 
in weighable amounts in pure form, still we know much about 
the chemical and physical properties of these substances and 
are thus able to learn their valence. For example, when we 
know that mesothorium is iso topic with radium and that the 
valence of the latter is two, we are safe in concluding that 
mesothorium has the same valence. Since UXj and ionium are 
isotopic with thorium we conclude that all three have the same 
valence, namely four; and since U 2 is isotopic with uranium itself 
its valence is six (836). These examples suffice to illustrate 
the principle employed. Let us now consider the first part of 
the uranium-radium series. In Table LV the first line gives the 

TABLE LV 





U 


UXx 


UX, 


U, 


Io 


Ra 


Em 


Rays 


a 

6 




4 




5 


a 

6 


a 

4 


a 

2 




Valence 


O 



symbols of the products of uranium in the order of their forma- 
tion (868); the next line shows the kind of rays given out by 
each substance; while the third line shows the valence of each. 
Fajans discovered a remarkable relation between the facts 
set forth in this table, namely, (1) the valence of the product of 
an alpha-ray change (858) is less by two than that of the parent; 
(2) the valence of the product of a beta-ray change (860) is greater 
by one than that of the parent. This statement is called Fajans 1 
Law. With a little amplification the law applies to all radio- 
active substances and their transformations. 



584 Introduction to General Chemistry 

873. The Meaning of the Term Element.— A direct defini- 
tion of the term element has not been given in this text (see 31). 
To say that an element is a substance incapable of further decom- 
position would exclude uranium, radium, and thorium as well 
as, of course, all other radioactive substances from the list of 
elements; and this would certainly be an unwise classification. 
Since radioactive changes are always spontaneous and entirely 
beyond human control, it would seem best to define the term 
element thus: An element is a substance which cannot be decom- 
posed into simpler substances at the will of the experimenter. We 
shall then class all radioactive products (even those like thorium 
emanation, period 53 seconds) as elements. All these "new" 
elements have been satisfactorily fitted into enlarged periodic 
tables. For example, see the lower part of the Harkins table 
(849). It will be seen that several elements occupy the same 
position. Except for the rare earths these are isotopes. 

874. The Technical Production of Radium. — Radium was 
first made in considerable quantities from pitchblende, a mineral 
having the composition U 3 Os. This mineral is so rare that it is 
no longer a satisfactory source of radium. The bulk of the 
radium now produced is made from carnotite, a mineral 
of somewhat variable composition, but containing uranium, 
vanadium, and potassium as its usual constituents. This ore 
is found chiefly in southwestern Colorado and adjoining regions 
of Utah, for the most part as cementing material in sandstone. 
The ordinary ore as mined carries less than 1 per cent of uranium 
and therefore but one part of radium in 300 million parts of 
weight of ore! There are several ways of extracting the radium, 
but all of them are based on the fact that the radium present 
remains with the barium present when the latter is separated. 
If insufficient barium is present in the ore, enough barium 
chloride is added to bring the proportion up to about 1 per cent. 
Usually the barium (with the radium) is first obtained as sulfate 
(164). This salt is converted into chloride and the latter 
subjected to fractional crystallization in the manner already 
described (856). A much more rapid separation of barium and 
radium results from the crystallization of the bromides, since 



Radioactivity and the Nature of Matter 585 

in this case above 95 per cent of the radium separates with half 
of the barium salt. In another process the hydroxides of barium 
and radium are crystallized, in which case the radium concen- 
trates in the mother-liquors instead of in the crystals. 

The world's total production of radium up to the end of 1918 
was between 100 and 125 grams, over half of which was made 
in the United States of America. 

875. How Radium Is Measured and Sold. — Radium usually 
comes on the market as an isomorphous (810) mixture of its 
bromide with about an equal amount of barium bromide. The 
selling price of radium is always based on the radium content of 
material, stated in milligrams of radium element. The radium 
content of a given sample is not found by the use of a balance 
but is determined by the ionizing power of the gamma rays (861) 
as shown by an electroscope (770, 854). Comparative measure- 
ments are always made with a tube of radium salt of exactly 
known radium content (a standard). Every purchaser of 
radium can, for a small fee, have the material measured by the 
United States Bureau of Standards and get a certificate before 
he pays for it. During the year 1918 the price was close to 
$120.00 per milligram of radium element. This is more than 
100,000 times the value of gold! But when it is known that 
a ton of ordinary ore yields only two milligrams of radium and 
that it requires three to four months' work to extract and refine 
the radium, the high price is readily understood. 

876. The Use of Radium in Therapeutics. — The term thera- 
peutics is defined as being that branch of medical science that 
treats of the action of remedial agents on the human body. The 
beta and gamma rays of radium are capable of producing serious 
" burns" of the skin and underlying tissues. The effect produced 
is proportional to the quantity of radium and the length of time 
of the exposure. Healthy tissue is much less affected by these 
rays than is abnormal and unhealthy tissue. It is by reason of 
this fact that radium finds an important use as a therapeutic 
agent. Exposure of certain kinds of abnormal growths on the 
body to radium rays results in their destruction and subsequent 
removal. Some forms of cancer respond very favorably to such 



586 Introduction to General Chemistry 

treatment. In cases where large masses of cancerous tissue are 
removed by surgical operation treatment with radium rays is 
often subsequently employed to destroy the remaining portions 
not accessible to the knife. In hopeless cases, where a cure 
cannot be expected, radium treatment is useful in greatly 
alleviating pain. In most respects radium rays produce thera- 
peutic effects like those due to X-rays. The great advantage 
over X-rays in the use of radium arises from the fact that the 
radium container is so minute that it can be applied exactly 
where its rays are required to act, a procedure impossible with an 
X-ray bulb. Radium in considerable quantity (100 mg. or more) 
may be safely handled if contained in thick-walled lead tubes. 

877. X-Ray Spectra. — Ordinary visible light, which is made 
up of ether vibrations of various wave-lengths (788), can be 
spread out into a spectrum of its component colors in another 
way besides that by the use of a prism, namely by means of a 
diffraction grating. The latter consists of a polished plane 
surface, as of glass or metal, ruled with an enormous number of 
fine parallel lines, often as little as one- thousandth of a milli- 
meter apart. The principle of the production of spectra by 
such gratings is discussed in most textbooks on physics. 

For many years after the discovery of X-rays numerous 
attempts were made to obtain X-ray spectra, but all without any 
success, by the use of either prisms or gratings. It then occurred 
to Laue that if X-rays consisted of very much shorter waves than 
visible light it would be impossible to rule gratings with lines 
sufficiently fine and close together to show the expected effect. 
In fact, he calculated that if X-rays consist of waves one- 
thousandth as long as those of visible light, in order to get their 
spectrum it would require a grating with lines no farther apart 
than the diameters of ordinary atoms (196). To rule such a 
grating would be a task beyond human skill. Then came the 
brilliant idea that nature makes such gratings on every hand, 
since in every crystal the molecules are arranged in rows and 
layers like bricks in a wall (204) . Several experimenters applied 
Laue's ideas and soon brought forth a great number of new and 
interesting facts, a few of which we shall discuss briefly. 



Radioactivity and the Nature of Matter 



587 



When X-rays from a given source strike the flat surface of 
a crystal obliquely the reflected rays, brought to focus on a 
photographic plate, produce a well-defined line spectrum con- 
sisting of a small number of lines. Each line corresponds to a 
definite wave-length (just as in the case of visible light), which 
length can be easily calculated. 

878. Atomic Numbers. — The X-ray spectrum given by one 
element used as the target of the X-ray bulb is characteristic for 
that element and different from that of any other element. In this 
connection a remarkably simple relation was discovered in 19 14 
by the young English physicist Moseley: the wave-length of the 



TABLE LVI 

Periodic Table According to Atomic Numbers, H:i 



Pe- 
riod 





I 
A B 


II 

A B 


Ill 

A B 


IV 

A B 


V 
A B 


VI 
A B 


VII 
A B 


VIII 


I 


2 
He 


3 
Li 


4 
Be 


5 
B 


6 
C 


7 

N 


8 



9 
F 




2 


10 

Ne 


11 

Na 


12 
Mg 


13 
Al 


14 
Si 


15 
P 


16 

s 


17 
CI 






18 
A 


19 
K 


20 
Ca 


21 22 

Sc Ti 


23 
V 


24 

Cr 


25 
Mn 


26 27 28 
Fe Co Ni 


3 




29 
Cu 


30 
Zn 


31 I 32 
Ga Ge 


33 

As 


34 
Se 


35 
Br 






36 
Kr 


37 
Rb 


38 
Sr 


39 40 

Y jZr 


41 
Cb 


42 
Mo 


43 

? 


44 45 46 
Ru Rh Pd 


4 




47 48 

Ag Cd 


49 
In 


50 
Sn 


51 
Sb 


52 
Te 


53 
I 






54 
Xe 


55 
Cs 


56 
Ba 


57 
La 


58 
Ce 


59 
Pr 


60 
Nd 


61 

? 




i> 




62 
Sa 


63 
Eu 


64 
Gd 


65 
Tb 


66 
Dy 


67 
Ho 


68 
Er 




6 




69. 
Tm 


70 
? 


71 
Yb 


72 
Lu 


73 
Ta 


74 
W 


75 
? 


76 77 78 
Os Ir Pt 




79 1 80 
Au Hg 


81 82 
Tl Pb 


83 
Bi 


84 

? 


85 

? 




7 


86 

Nt 


87 88 
? Ra 


89 90 

? Th 


9i 

? 


92 



588 Introduction to General Chemistry 

most intense line decreases quite uniformly as we pass from one 
element to the next higher in atomic weight. Moseley also 
showed that if the elements are assigned numbers, called 
atomic numbers, in the order of their increasing atomic weights 
there is a very simple numerical relation between the atomic 
number of an element and the wave-length of the strongest line 
of its X-ray spectrum. The atomic numbers calculated from 
the wave-lengths as found by experiment differed from whole 
numbers by not more than 1 or 2 per cent, and this difference 
could easily be ascribed to error of experiment. The only cases 
where the atomic numbers so found did not follow exactly the 
order of increasing atomic weights were met with in those pairs 
of elements (A and K, Te and I) where a departure from the 
usual order is necessary to bring the elements into their proper 
groups in the Periodic Table (829). In Table LVI the elements 
are arranged in a periodic table strictly in the order of their 
atomic numbers. The anomalies (848) now disappear com- 
pletely. It would therefore appear that the atomic number of 
an element is its most fundamental constant. In Table LVI 
the rare earth elements (820, 846) are included within the heavy 
lines. These elements form a single family of Group III (849) 
and do not belong to the other groups indicated by their 
positions in this table. 

879. The Structure of Crystals. — Professor W. H. Bragg and 
his son W. L. Bragg, English physicists, have followed Laue's 
ideas in another direction and obtained a great deal of informa- 
tion regarding the arrangement of molecules and atoms in a 
great variety of crystals. The simplest way of investigating the 
structure of a crystal consists in getting the impression made on 
a photographic plate when a small, round beam of X-rays passes 
through a thin slice of the crystal and strikes the plate some 
inches beyond. The reproduction of such a photograph is shown 
in Fig. 121. From the nature of the pattern produced the geo- 
metrical arrangement and distances from one another of the 
atoms of a crystal can be calculated. 

880. Atomic Structure and the Nature of Matter. — The 
modern view of the structure of an atom has been described 



Radioactivity and the Nature of Matter 



589 




briefly in section 470. The facts epitomized in the periodic law, 
together with the phenomena of radioactivity, show very con- 
vincingly that atoms must be constructed according to very 
definite plans, and that the arrangement of the parts (electrons, 
atoms of helium, and the 
atomic nucleus, 470) in the 
atom of an element must de- 
termine its properties. Some 
of the world's ablest physi- 
cists and chemists have de- 
voted much attention to the 
problems of atomic structure, 
and as a result several more 
or less definite hypotheses 
have resulted, each aiming to 
explain and correlate as many 
facts as possible. As yet none 
of these suggestions is wholly 
satisfactory, although several 

of them account very well for many of the facts with which 
we are familiar. 

On one point there is general agreement: namely, that the 
atomic number of an element represents the number of electrons 
encircling the electro-positive nucleus of the atom. Since atoms, as 
such, are electrically neutral, the positive charge of the nucleus 
must also be proportional to the atomic member. It seems prob- 
able that the nucleus also contains some electrons and that its 
positive charge represents the excess of its positive over its nega- 
tive electricity. 

In changes of atoms into ions and vice versa (oxidations and 
reductions, 501-507) the valence electrons lost or gained are those 
of the outer ring, and not of the nucleus. In radioactive changes 
the alpha and beta rays doubtless come from the nucleus, so that 
the new residual atom has a nucleus different in positive charge 
from its parent and therefore surrounded by a different number 
of electrons in the outer rings. It is on this basis that Fajans' 
Law (872) receives its explanation. 



59° Introduction to General Chemistry 

It is of interest to note that the chemist's control over the 
composition of the atom is limited to the removal and replace- 
ment of some of the electrons of the outer rings; changes of the 
nucleus are solely the result of spontaneous radioactive pro- 
cesses entirely beyond human control. 

Since the positive charge of the nucleus is the excess of its 
positive over its negative electricity, it is possible for two or more 
differently composed nuclei to have equal nuclear charges. 
Atoms containing nuclei differing thus would have equal 
numbers of electrons in the outer rings and therefore be identi- 
cal in chemical but not in radioactive properties. In conse- 
quence it would be possible to have two (or more) different 
elements with identical chemical properties, in other words 
isotopes (871). 

In conclusion, it may be stated that all chemists and 
physicists are now agreed that the elements are but modifications 
of the same primitive forms of material, including in the latter 
term electrons and whatever else may compose an atom. It 
may well be that all matter is entirely of an electrical nature. 
This view has led some persons to declare that "there is no 
such thing as matter," that " every thing is electricity." This 
statement is scarcely warranted. It is just as if, when we have 
discovered that a potato is composed of starch, fiber, and water, 
we should declare "there is no such thing as a potato"! As for 
matter, it remains what it always has been, only we know more 
about it. 



INDEX 



INDEX 

[References are to sections, not pages] 



Absolute, temperature, 6 

Absorption, 731 

Acetamide, 659 

Acetates, precipitation by sodium ace- 
tate, 452 

Aceteldehyde, 652 

Acetic acid: as a dissolving agent, 457; 
effect of acetates on ionization of, 431 ; 
effect of hydrochloric acid on, 432; 
ionization of, 409; failure to precipi- 
tate acetates, 452; graphic formula, 
654; general properties, 157, 653; 
and sodium hydroxide, 435; titra- 
tion of, 440 

Acetone, 656 

Acetylene: composition, 50; and copper 
oxide, 83; heat of combustion, 357; 
preparation, 49; series, 663 

Acheson process, 630 

Acid anhydride, 313 

Acidimetry, 137 

Acids: and carbonates, 163; dibasic, 
102; fatty, 655; ionization of, 410; 
monobasic, 102; parts of, 377; prop- 
erties of, 90; strength of, 428; 
tribasic, 159; strong: little soluble 
salts of, and acids, 458; salts of, and 
weak acids, 457; and salt of weak 
acid, 428, 430; titration of, 440; and 
weak acid, 432; and weak base, 437; 
weak, 177; little soluble salts of, solu- 
tion by strong acids, 456; as precipi- 
tating agent, 452 ; and strong acid on, 
432; and a strong base, 434; suppres- 
sion of ionization of, 431 ; titration of, 
440; and a weak base, 438 

Adhesion, 726 

Adsorption, 728, 731, 732, 763; accom- 
panying precipitation, 739; and 
catalysis by finely divided metals, 
731; by colloid, 739; of gases by 
charcoal, 728; from solution, 732; 
from solution, explanation of, 763 

Affinity, 259 

Agate, 804, 806 1 

Air: adsorption by glass, 729; compo- 
sition of, 10, 765, 766; ionization of, 
854; liquid, 776, 777; weight of, 3 

Aitken, 769 

Albumin, 685; and the phosphoric 
acids, 598 



Alcohol: absolute, 641 ; denatured, 641; 
from fermentation, 640, 641; wood, 
645 

Alcohols: aromatic, 670; triatomic, 679 

Aldehydes, 652; aromatic, 670 

Alfalfa, assimilation of nitrogen, 515 

Aliphatic compounds, 666 

Alkali metals, 823 

Alkaline earth metals, 824 

Allotropic forms, 582, 600 

Alloys, fusible, 812 

Alpha rays, 480, 858 

Alum, 175 

Aluminium. See Aluminum 

Aluminum, 174; acid reaction of salts 
of, 176; place in Periodic Table, 833; 
chloride, 174; hydroxide: gel, 761; 
precipitation by ammonium hydrox- 
• ide, 452; preparation and properties, 
174, 177; nitrate, 175; nitride, 514; 
oxide, 174; potassium sulfate, 175; 
sodium sulfate, 175; sulfate, 175' 

Amethyst, 806 

Amides, 659 

Amine acids, 685 

Amines, 658; aromatic, 674 

Ammeter, 400 

Ammonia: adsorption by charcoal, 728; 
composition of, 52, 53; and copper 
oxide, equation, 84; critical tempera- 
ture of, 775; liquid, 517; manufacture 
from cyanamid, 526; manufacture of 
synthetic, 525; oxidation to nitric 
acid, 570; properties of, 51, 517, 527; 
sources of, 516; theory of the syn- 
thesis of, 520-24; uses of, 518; and 
water, 91; and water, equilibrium, 
284. 

Ammonium: aluminum sulfate, 175; 
arseno-molybdate, 810; bicarbonate, 
dissociation of, 530; chloride, 92; dis- 
sociation of vapor, 529; and sodium 
hydroxide, ionic theory of, 426, 402; 
cyanate, synthesis of urea, 696; 
fluoride, 269; hydrogen sulfate, 101; 
hydroxide: a base, 91; degree of 
ionization of, 409; effect of am- 
monium salts on ionization, 431; 
effect of sodium hydroxide on ioniza- 
tion, 432; and nitric acid, 105; 
titration of, 440; molybdate, 507; 



593 



594 



Introduction to General Chemistry 



[References are to sections, not pages] 



nickel sulfate, 817; nitrate, 105; 
decomposition by heat, 565; use as 
an explosive, 556, 573; nitrite, 513; 
perchlorate, 355; phosphomolybdate, 
597,813; picrate,673; sulfantimonate, 
811; sulfate, 101; sulfides, 607; disso- 
ciation of, 530; yellow, 607; sulfo- 
cyanate and ferric chloride (ionic 
equilibrium), 280, 415 

Ampere, 400 

Amphoteric substances, 177 

Analysis, 31; by means of spectroscope, 
786 

Anaxagoras, 184 

Anhydride of an acid, 313 

Aniline, 674 

Animals, dependence on plants, 690 

Anions, 389, 391 

Anode, 295, 389 

Anthracene, 698 

Antichlor, 611 

Antimony: and its compounds, 811; 
place in Periodic Table, 835; tri- 
chloride, from antimony and chlorine, 
246 

Apatite, 580 

Aqua regia, 562 

Aquadag, 750 

Argol, 665 

Argon, 513; in air, 765; discovery and 
properties, 791; family, 825; molec- 
ular and atomic weight, 792; place 
in Periodic Table, 848 

Argyrol, 750 

Aromatic series, 666 

Arrhenius, Svante, 405, 720 

Arsenic: antidote for, 810; and com- 
pounds, 810; place in Periodic 
Table, 835 

Arsenious sulfide, 810; colloidal sus- 
pension, 734, 736-38 

Arsine, 810 

Atmosphere, 765-99; a disperse system, 

774 

Atomic-molecular hypothesis, 208 

Atomic numbers, 848, 878, 880. See 
flyleaf at back of book 

Atomic structure, 880 

Atomic volume, 823 

Atomic weights, 216; determined by 
law of Dulong and Petit, 229; oxygen 
basis, 223; of radioactive substances, 
869; relation to symbol weight, 216; 
table, 800. See also inside of back 
cover of book 

Atoms: Greek conception of, 184; mod- 
ern conception of, 185; number in a 
molecule, 214, 215; relative weights 
of, 208-13; structure of, 470, 880 



Attraction, molecular, 202, 726 

Atwater, 687 

Avogadro's hypothesis (or law), 193; 
application to suspensions in liquids, 
707; applications, 210; exactness of, 
226 

Avogadro-Van't Hoff hypothesis, 716 

Azote, 512; (nitrogen) 15, 

Babbitt metal, 811 

Baking powder: phosphate, 593; tar- 
trate, 665 

Balloons: helium, 797; hydrogen, 304 

Barium, 824; flame, 784; in Periodic 
Table, 832; chloride, 164; peroxide, 
310, 319; salts, reaction with sul- 
fates, 380; sulfate, 164 

Base: diacid, 146; ionization of, 410; 
monacid, 146; parts of, 378; strength 
of, 429; strong: and salt of weak 
base, 429-30; titration of, 440; and 
weak base, 432; weak, 176; and 
strong acids, 437; and strong base, 
43 1 ; suppression of ionization of, 43 1 ; 
titration of, 440; and weak acid, 438 

Basic nitrates, 564 

Batteries, electric, 496 

Baume, scale of specific gravity, 618 

Be, 618 

Bead test, metaphosphate, 596; borax, 
803 

Becquerel, 852; rays, 853 

Beer, 640 

Beets, sugar, 682 

Bell metal, 809 

Benedict, 687 

Benzaldehyde, 670, 697 

Benzene, 667; ring, 668 

Benzene sulfonic acid, 672, 697 

Benzoic acid, 671 

Benzyl alcohol, 670 

Beri-beri, 689 

Beryllium, 826; in Periodic Table, 832 

Beta rays, 860 

Birkland-Eyde process, manufacture of 
nitric acid, 568 

Bismuth: compounds, 812; place in 
Periodic Table, 835; nitrate, 564; 
subnitrate, 564 

Bivalent, 146 

Bleaching, by chlorine, 249 

Bleaching powder, 351 

Blotting paper, functioning of, 726 

Blowpipe, oxyhydrogen, 300 

Bluestone, 497 

Bedy: living, need of food, 686; use of 
word in chemistry, 22 

Boiling-point: molar elevation, 718; of 
solutions, 128; of water, 112 



Index 



595 



[References are to sections, not pages] 



Boisbaudran, 845 

Bonds: double, 661; triple, 663; va- 
lence, 323 

Bone ash, 158 

Bone black, use as adsorption agent, 732 

Boracic acid, 801 

Borax, 801, 803 

Boric acid, 801, 802 

Boron: flame, 784; in Periodic Table, 
833; properties, 801 

Boyle's Law, 4; explanation of, 188 

Bragg, W. H., 879 

Bragg, W. L., 879 

Brandt, 577 

Brass, 148 

Breakfast, example, 688 

Bredig, 744 

Brin's process, for oxygen, 310 

Britannia metal, 811 

British Thermal Unit, 358 

Bromic acid, 822 

Bromides: action of chlorine on, 259; 
insoluble, 257 

Bromine: liquid, 255; occurrence of, 
254; in Periodic Table, 83 7 ; prepara- 
tion and properties, 255, 258; uses 
of, 260; water, 255 

Bronze, 809 

Brownian movements: of particles in 
liquids, 705; of smoke particles, 703 

Brownlee apparatus, 44 

B.T.U., 358 

Bunsen burner, 781 

Burning of substances, 10-20 

Burns: from nitric acid, 104; by steam, 
117; from sulfuric acid, 93 

Butyric acid, 680 

Butyrin, 680 

Cadmium, 832 

Caesium, 823; most active metal, 842; 
in Periodic Table, 831 

Calcium, 150, 832; flame, 784'; spec- 
trum of, 786; bicarbonate, 156; 
carbide, 49, 631; carbonate, 150; 
dissolved by hydrochloric acid, 456, 
461; precipitation of, 448; chloride, 
151; action with carbonic acid, 449; 
cyanamid, 514; fluoride, 267; hy- 
droxide, 151, 166; hypochlorite, 351; 
nitride, 514; oxide, 150; phosphate, 
158; precipitation of, 452; silicate, 
270; sulfate, 153; tungstate, 814 

Calomel, 182 

Calorie, in 

Calorimeter, bomb, 357 

Candle, burning of, 20, 780 

Cane sugar, 682 

Caprilic acid, 680 



Caproic acid, 680 

Carat, 819 

Carbides, 631 

Carbohydrates, 682 

Carbolic acid, 672 

Carbon: compounds, 629-701; "fixed," 
in fuel, 359; formation in flame, 782; 
heat of combustion, 357; and hydro- 
gen, 631 ; place in Periodic Table, 826, 
834; properties of, 630, 631; as 
reducing agent, 328; valence, 648; 
weight in one liter of gaseous com- 
pounds, 60; bisulfide (or disulfide), 
546 and 631; dioxide, 19, 30, 633; 
adsorption by charcoal, 728; from 
calcium carbonate, 150; composition 
of, 39; constancy of concentration in 
air, 766; critical temperature of, 775; 
and limewater, 151; solid, as a refri- 
gerant, 633, 775; and water, equilib- 
rium, 285; disulfide, 631; monoxide, 
632, 665; critical temperature of, 775; 
heat of combustion, 357; liquefaction 
of, 777; reducing agent, 329; oxy- 
chloride, 695; tetrachloride, 644 

Carbonates and acids, 163 

Carbonic acid, 152; action with calcium 
chloride, 449; equilibrium, carbon 
dioxide, and water, 285; ionization 
of, 409; and sodium hydroxide, 161 

Carborundum, 631 

Carnotite, 874 

Caro's acid, 628 

Catalysis by metals and adsorption, 731 

Catalytic agent, or catalyzer, 239; for 
ammonia reaction, 522; for decom- 
position of hydrogen peroxide, 320; 
manganese dioxide in preparation of 
oxygen, 306; manufacture of sul- 
furic acid, 615; platinum, 303 

Cathion, 389, 390 

Cathode, 295, 389; rays, 475, 478 

Caustic soda, absorbs carbon dioxide, 
19. See Sodium hydroxide 

Cavendish, 292, 544 

Cells, galvanic, 496 

Celluloid, 694 

Cellulose, 684; products, 694 

Cerium, 820, 834 

Cerium oxide, use of, in gas mantles, 301 

Chamber acid, 616 

Chamber process, 616 

Chaptal, 572 

Charcoal: adsorption by, ;jS, 732; 
burning of, 18 

Charge on a suspension, test for, 749 

Charles's Law, 5 

Chemical activity of elements, order of, 
in Periodic Table. 842 



596 



Introduction to General Chemistry 



[References are to sections, not pages] 



Chemical reactions, rate of, 275-78 

Chile saltpeter, 104; iodine from, 261 

Chlorates, preparation of, 353 

Chloric acid, 354; and lead sulfide, 354 

Chlorides, insoluble, 252 

Chlorination process for gold, 819 

Chlorine: and aluminum, 174; and 
bromides, 259; critical temperature, 
775; discovery of, 234; from elec- 
trolysis of hydrochloric acid, 43; 
electrolytic preparation, 237, 238; 
and ferrous chloride, 173; and hydro- 
gen, 44, 243; liquefaction of, 242; 
and mercury, 179; and methane, 644; 
minimum weight, 63; occurrence, 
233; oxidation products of, 352; an 
oxidizing agent, 332; in the Periodic 
Table, 827; and phosphorus, 247; 
physical property of , 241; poisonous 
gas, 236; preparation from hydro- 
chloric acid, 235, 239; and sodium, 
theory of union of, 485; and tur- 
pentine, 248 ; union with metals, 246 ; 
and uses of, 249; and water, at 
ordinary temperatures, 245; at high 
temperatures, 240; dioxide, 354; 
monoxide, 352 

Chloroform, 644 

Chlorophyl, 691 

Chloropicrin, 695 

Chromates, 345; as oxidizing agents, 
346 

Chromium: and its compounds, 344; 
place in the Periodic Table, 836; 
hydroxide, precipitation of, 452 

Cinnamic aldehyde, 697 

Cinnamon, oil of, 697 

Citric acid, 665 

Claude, 776 

Clay, 177, 804 

Cleveite, helium from, 794, 867 

Clouds, 768-71; functions of gaseous 
ions in their formation, 771 

Clover, assimilation of nitrogen, 515 

Cloves, oil of, 697 

Coal, distillation of, 634 

Coal tar, 667 

Cobalt: and its compounds, 816, 817; 
place in Periodic Table, 838 

Coefficient, in chemical equations, 76 

Cohesion, 726 

Coke, 630; as a reducing agent, 328 

Collodion, 694 

Colloids, 735; adsorption of agent 
which precipitates, 739; chemistry, 
importance of, 764; protecting agent 
for suspensions, 745; test for charge 
on, 749 



Columbium, 835 

Combining volumes, Gay Lussac's Law, 
220, 221 

Combustion: heat of, 356-58; in nitric 
oxide, 546; spontaneous, 364. See 
also Burning of substances 

Common ion law, 432 

Complex ions, 538 

Composition, law of definite, 46, 99 

Compound, 31 

Concentration and speed of reaction, 
280 

Conductivity: change during neutraliza- 
tion, 423; effect of dilution, 406; 
molecular, 407 

Conservation of energy: for bodily 
processes, 687; law of, 371 

Conservation of matter, 21 

Constant heat summation, law of, 363 

Constant of equilibrium (example), 283 

Constant proportion, law of, 46, 99 

Contact process for sulfuric acid, 617 

Cooledge, 814 

Copper, 165; bead test for, 783; burning 
of, 32; flame, 784; from hydrogen 
and copper oxide, 33 ; and nitric acid, 
550, 561; properties and place in 
Periodic Table, 831 {see also Cupric 
and Cuprous); ammonium ion, 538; 
chloride, 165; from chlorine and 
copper, 246; hydroxide, 165; nitrate, 
165; oxide, 165; and acetylene, equa- 
tion, 83; and ammonia, equation, 84; 
composition of, 32, 38; and hydrogen, 
33, 82; sulfate, 165 

Cordite, 693 

Corn sirup, 639 

Corrosive sublimate, 178 

Cotton, 684; bleaching of, 351; soluble, 
694 

Coulomb, 400 

Cream of tartar, 665 

Critical pressure, 775; temperature, 775 

Crookes, 477, 859; tube, 475 

Cryolite, 267 

Crystals: melting of, 205; theory of 
structure, 204; X-ray and structure 
of, 879 

Cupric compounds {see also Copper 
compounds), 333; bromide, color 
of, in solution, 396; chloride, elec- 
trolysis of solution, 386; as catalyst, 
239; oxide, 325; potassium chloride, 
175; sulfate, electrolysis of solution, 
386; sulfide, precipitated by hydro- 
gen sulfide, 452 

Cuprous compounds, 2,33', oxide, 325 

Curie, Mme, 855 



Index 



597 



[References are to sections, not pages] 



Current, electrical : by chemical action, 
493-95; direction of, 472; electron 
theory, 469; how carried through a 
wire, 471; nature of, 468; strength 
of, 400 

Cyanates, 665 

Cyanide process for gold, 819 

Cyanides, 665 

Dalton: Atomic Hypothesis, 208; Law 

of Partial Pressure, 192 
Daniell cell, 496 
Deacon's process, 239 
Decomposition, 25; of sal soda, 26 
Degree of ionization, 408-10 
Deliquescence, 130 
Democritus, 184 
Density, no 
Depilatory, 607 
Detonator, 573 
Developers, photographic, manufacture 

related to dye industry, 699 
Dew, 767 

Dewar, 777, 779; vessels, 777 
Dew point, 767 
Dextrose, 639 
Dialysis, 743 
Dialyzer, 743 

Diamonds, 630; and gamma rays, 861 
Diastase, 682 

Diatoms (infusorial earth), 732 
Dichlor diethyl sulfide (mustard gas), 

6 95 ' 

Dichromates, 345; oxidizing agents, 346 

Dietetics, 688 

Diffusion: of gases, 191; in liquids, 704 

Dinner, example, 688 

Disintegration hypothesis, 481, 865 

Disperse systems, 725-64, 774 

Displacement: electronic interpreta- 
tion of, 489; metallic, electronic 
interpretation of, 491; of metals by 
one another, 490; of non-metals by 
one another, 488 

Dissociation: electrolytic {see Ioniza- 
tion); hydrolytic, 436; of volatilized 
ammonium salts, 529-31 

Distillation, 23 

Dithionic acid, 628 

Double-decomposition, 337, 383; and 
electrical conductivity, 384; and the 
ionic hypothesis, 413 

Drying agents, 130 

Dulong and Petit, law of, 229-30 

Durion, 539, 804. 

Dust: in the air, 768; counting of 
particles, 769; explosions, 365; func- 
tion of, in cloud formation, 768 

Dynamite, 692; preparation of, 726 



Effervescence, 163 

Efflorescence, 131 

Egg white, 685 

Ekaboron, 845 

Ekaluminium, 845 

Ekamanganese, 845 

Ekasilicon, 845 

Electric cells, oxidation-reduction, 502 

Electric current, 468, 469; oxidation 
and reduction by means of, 507 

Electricity: frictional,474; nature of, 466 

Electrochemical equivalents, law of, 403 

Electrodes, potential difference of, 499 

Electrolysis, 27; of cupric chloride 
solution, 386; of cupric sulfate solu- 
tion, 386; Faraday's laws of, 399; 
of hydrochloric acid, 385; of silver* 
nitrate, 387; of sodium chloride solu- 
tion, 385; terms used in, 389; theory 
of, 398, 487; of water, 27 

Electrolytes, 389; molecular weights 
of, 720; precipitation of, by common 
ion, 453; soluble, equilibrium be- 
tween, 441; solution of little soluble, 

455 

Electromotive force, 499 

Electromotive series of metals, 492, 499 

Electrons, 465, 466; mass of, 479; 
proof of existence, 467; vibration, 
cause of light, 788 

Element, 31, 35; meaning of term, 
873; total number, 800 

Elements, activity of, 842 

Emery, 174; powder graded by time 
of settling in water, 733 

Emulsifying agents, 755 

Emulsoids, 735, 751-57; general prop- 
erties, 756; importance, 757 

Energy: chemical, 372; chemical, con- 
version into electrical energy, 50S; 
conversion of, 371; electrical, 500; 
electrical, conversion into chemical 
energy, 509; forms of, 372; kinetic, 
of molecules, 198 

Enthodermic changes, 366 

Enzymes, 682 

Epsom salts, 144 

Equations, chemical, 76; balancing of, 
86; balancing of oxidation and reduc- 
tion, 561; meaning of, 77, 85; prob- 
lems of, 79, 87; review table, So 

Equilibrium (general): chemical. 274; 
constant (example), 2S3; criterion of, 
282; effect of changes of concentra- 
tion on, 280; effect of pressure on 
hydrogen and nitrogen, 523; effect 
of pressure on system in, 287; effect 
of removing one product of the re- 
action, 2S9; between electrolytes {see 



598 



Introduction to General Chemistry 



[References are to sections, not pages] 



Equilibrium betweeen electrolytes); 
and heat production, 367; hydrogen 
and nitrogen, effect of temperature 
on, 288, 521; kinetic hypothesis 
applied to, 279; liquid and vapor, 
201 ; between molecules and ions, 405 ; 
physical, 273 

Equilibrium between electrolytes: and 
gas evolution, 459; graphic repre- 
sentation of, 417, 418 ff., 444; be- 
tween soluble electrolytes, 441; in 
solution, 413-39; in solution and solid 
substance, 443 ff. 

Equivalent, electrochemical, 403 

Equivalent weight, 403 

Esters, 657; of glycerine, 679 

Ethane, 643 

Ether, 642 

Ethyl compounds: acetate, 657; alco- 
hol, 641; structural formula, 649; 
ammonium iodide, 658; iodide, 660 

Ethylene, 660; series, 662; structural 
formula of, 661 

Ethylene chloride, 660 

Evaporation, theory of, 199 

Exothermic changes, 366 

Explosion, 302 

Explosives, 571, 692, 693 

Fajans' Law, 872 

Falk, 722 

Families: of elements, 822; A and B, 831 

Faraday, 242, 389, 403, 775; law of 
electrolysis, 399 

Fat soluble A, 689 

Fats, 677; composition of, 680 

Fehling's solution, 683 

Fermentation, 640 

Ferric: chloride, and ammonium sul- 
focyanate, 280, 415; and hydrogen 
sulfide, 503; and potassium iodide, 
503; preparation, 173; compounds, 
331; hydroxide, 173; colloidal, 741- 
43; a gel, 761; and hydrochloric 
acid, 455; precipitation of, 451; 
oxide, 173; sulfate, 173; and sodium 
carbonate, 384 

Ferro-molybdenum, 813 

Ferro-silicon, 804 

Ferro-tungsten, 814 

Ferrous compounds, 331; ammonium 
sulfate, 175; chloride, 173; hy- 
droxide, 173; oxide, 173; sulfate, 
173; sulfide: dissolved by hydro- 
chloric acid, 456; precipitation of, 
452; preparation and use, 339, 601 

Fertilizer: phosphate, 160; use of 
ammonium salts in, 518 



Films, photographic, 694 

Filters, adsorption by, 763 

Filtration, 23 

Fire extinguisher, 633, 644 

Firefly, 584 

Flame, 780; colored, 784; reactions in, 
783; spectra, 786 

Flint, 806 

Flour, bleaching of, 563 

Fluorides, 269 

Fluorine: most active non-metal, 842; 
in Periodic Table, 837; preparation 
and properties, 267 

Fluor-spar, 267 

Fluosilicates, 272 

Food, 676 

Formaldehyde, 652 

Formalin, 652 

Formic acid, 665 

Formula: calculation of, 80, 81; chemi- 
cal, 62; of elementary gases, 75, 218; 
graphic, 323; involatile substances, 
72; making of, 67; structural, of ethyl 
alcohol, 649; structural, importance 
of, in organic chemistry, 651; struc- 
tural, of methyl ether, 649; use of, 68; 
volatile, liquids and solids, 7 1 ; weight, 
relation to molecular weight, 217; 
weight, and symbol weight, 74 

Fox fire, 584 

Frasch process, 602 

Fraunhofer lines, 790 

Frazier, 712 

Freezing-point, molar depression, 718 

Frost, 767 

Fruit sugar, 682 

Fuel: composition of, calorific power, 
359; for steam production, 360 

Fullers earth, 732 

Gallium, discovery of, 845 

Galvanic cells, 498 

Galvanized iron, 148 

Gamboge, Brownian movements of, 705 

Gamma rays, 861 

Gas: calorific power, 359; collected over 
water, calculation of pressure, 113; 
diffusion, 191; evolution of, factors 
governing, 463 ; evolution of, and ionic 
equilibrium, 459; illuminating, 634; 
natural, 643; pressure, cause of, 188; 
pressure, law of partial, 192; statistics, 
194-97 

Gas laws: accuracy of, 225; applica- 
tion to dilute solutions, 715; prob- 
lems on, 7 

Gases: the inert, or noble, 825; ioniza- 
tion of, 770; liquefaction of, 775; 



J 



Index 



599 



[References are to sections, not pages] 



mixing of, 190; standard conditions 
for, 7 

Gasoline, 643 

Gay Lussac's Law, 5 

Gelatine: gel, 758; a protecting agent 
for colloids, 745 

Gels, 758-61; plant and animal tissue, 
relation to, 759 

German silver, 816 

Germanium: discovery of, 845; place 
in Periodic Table, 834 

Glass: adhesion of water to, 726; 
adsorption of air on surface, 729; 
adsorption of water vapor on, 730; 
composition, 270, 808; etching of, 
271; optical, 803; quartz, 808; vari- 
ous kinds, 808 

Glucose, 639; fermentation of, 640 

Glycerine, 679 

Gofd: colloidal, 746; compounds and 
alloys, 819; extraction from ores, 819; 
place in Periodic Table, 831; solu- 
tion in aqua regia, 562; world- 
production of, 819 

Goldschmidt, 330 

Graham, Thomas, 735 

Grain alcohol, 641 

Granite, 804 

Grape sugar, 639 

Graphite, 630 

Gravity batteries, 497 

Grindstone, 806 

Groups: of atoms in organic chemistry, 
651; of elements, 830-38 

Guncotton, 693 

Gunpowder, 326, 571, 572 

Gypsum, 155 

Haemoglobin and oxygen, 314 

Halogens, 231, 822 

Hampson, 776 

Harkins, 754, 849 

Hartshorn, 516 

Heat: atomic, 230; of combustion, 
356-58; of formation, 361; of ioniza- 
tion, 439; latent, of evaporation, 115; 
latent, fusion of ice, 118; law of con- 
stant heat summation, 363; mechani- 
cal equivalent of, 369; molecular, 
792-93; of neutralization, 362, 439; 
production and equilibrium, 367; 
production in physical and chemical 
changes, summary of, 366; of reac- 
tion, 361; of solution, 127; of solu- 
tion, 288; of solution and solubility, 
134, 288; theory of, 189 

Heliotrope, 697 



Helium: and the alpha rays, 795, 858; 
balloons, 797; discovery of, 794; and 
family, 825; properties of, 796 

Hematite, 328 

Henry, Law of, 126 

Hess, Law of, 363 

Hexathionic acid, 628 

Hexoses, 682 

Hillebrand, 794 

Humidity of air, 766 

Hydrates, 97 

Hydrazine, 531 

Hydriodic acid, 265, 339 

Hydrobromic acid: oxidation of, 258; 
preparation of, 256; as reducing 
agent, 341; and silver nitrate, 257 

Hydrocarbons, 643; aromatic, 668; 
isomerism of, 664 

Hydrochloric acid: and aluminum, 174; 
and aluminum hydroxide, 175; and 
calcium carbonate, 461; and caustic 
soda, 41; and copper oxide, 165; 
electrolysis of, 43, 385; and iron, 173; 
and lead, 167; and lead dioxide, 157; 
and magnesium, 149; and magnesium 
hydroxide, 143, 455; and magnesium 
oxide, 145; preparation, 103, 250; 
properties, 251, 252; as reducing 
agent, 341; and sodium acetate 
(ionic theory of action), 424; and 
sodium hydrogen sulfate, 253, 289; 
and zinc, 149 

Hydrocyanic acid, 665 

Hydrofluoric acid, 269 

Hydrofluosilicic acid, 269, 272 

Hydrogen: adsorption by charcoal, 728; 
critical temperature, 775; discovery 
of, 292; from electrolysis of water, 
27; flame, temperature of, 299; heat 
of combustion, 357; liquid, 779; in 
nature, 292; percentage in water, 36; 
place in electromotive series, 492, 
499; place in Periodic Table, 847; 
preparation, 27, 28,33,293-95; prop- 
erties of, 27, 296, 297; use of, 295, 
304; weight in one liter of gaseous 
compounds, 56; weight in 22.4 liters 
of gaseous compounds, 56, 63. 211; 
reactions of: burning in chlorine, 
244; and chlorine, 44; and copper 
oxide, 33, 82; and iodine, chemical 
equilibrium, 281; and iodine, heat 
of reaction, 367; and iron oxide, 
290; from magnesium and steam, 
28; and nitrogen, 298; and nitro- 
gen, equilibrium, 521, 523; and 
oxygen, 299-303; chloride, 44; com- 
position, 48; critical temperature of. 



6oo 



Introduction to General Chemistry 



[References are to sections, not pages] 



775; physical properties of, 251 {see 
Hydrochloric acid); preparation of, 
44; fluoride, 268; action on quartz, 
silicon, glass, 270; iodide: prepara- 
tion, properties, 264; sulfide: aqueous, 
as reducing agent, 609; precipitation 
of sulfides, 452; preparation of, prop- 
erties, 339, 605-9; peroxide, 318-20; 
detection of (reaction with chromic 
acid), 321; as oxidizing agent, 347; 
as reducing agent, 348; use of, 320 

Hydrolysis of salts, 436 

Hydrosulfurous acid, 606 

Hydroxides, preparation of insoluble, 
166 

Hydroxylamine, 531 

Hypo, 625, 627 

Hypochlorites, 350 

Hypochlorous acid, 349; from chlorine 
monoxide, 352 

Hyposulfurous acid, 628 

Ice: density of, 119; latent heat, of, 
118; manufacture of artificial, 519 

"Icy-Hot" bottles, 777 

Ignition temperature, 302 

Illuminating gas, 634 

Indestructibility of matter, law of, 21 

Indicators, 440 

Indigo, synthesis of, 698 

Infusorial earth, 732 

Ink: sympathetic, 817; India, 750 

Insulator, 473 

Intumescence, 803 

Iodic acid, 822; and sulfurous acid, 277 

Iodides, 265; uses of, 266 

Iodine, 261-65; heat of reaction with 
hydrogen, 367; and hydrogen, chemi- 
cal equilibrium, 281-83; place in 
Periodic Table, 848; and starch, 263, 
637; uses of, 266 

Ionic equilibrium, 405, 413 ft.; and 
gas evolution, factors governing, 463 

Ionic hypothesis, 411; criticism of, 412; 
value of, 464 

Ionization: cause of, 486; degree of, 
conductivity method, 408-10; degree 
of, freezing-point method, 721, 722; 
degree of, graphic representation, 
417; of gases, 770; heat of, 439 

Ions: in solution: charges on, 404; and 
chemical reactions, 392; complex, 538; 
color of, 395; migration of, 397; 
nature of, 483 ; positive and negative, 
393; union of, 394; gaseous, 771 

Iridium and its compounds, 818, 838 

Iron, 173; burning of, 12, 17, 29, 81; 
galvanized, 148; and hydrochloric 



acid, 173; magnetic oxide of, 173; 
place in Periodic Table, 838; prepara- 
tion of, from hematite, 328; and 
steam, 29, 290; and sulfuric acid 
(concentrated), 621; and sulfuric acid 
(dilute), 173 

Iron compounds {see Ferrous and 
Ferric): iron oxide, 173 

Isatine, 698 

Isomerism, 647, 650 

Isomorphism, 810 

Isoprene, 700 

Isotopes, 871, 873 

Joule, mechanical equivalent of heat, 
370; unit of electrical energy, 500 

Kerosene, 643 

Ketones, 656 

Kieselguhr, 732 

Kinetic energy, 368; of molecules, 198 

Kinetic theory, 187; of the liquid state, 

198, 704 
Kipp, apparatus, 294 
Kraft, 577 
Krypton, 798; and its family, 825 

Lactic acid, 665 

Lactose, 682 

Langmuir, 754 

Lanthanum, 820; in Periodic Table, 
833, 846 

Latent heat: of evaporation, 115; of 
fusion, 118 

Laughing gas, 556 

Laurie acid, 680 

Lavoisier, 13 

Lead, 167; place in Periodic Table, 834; 
•acetate, 167; chloride, 167; dioxide, 
167; oxidizing agent, 326;fluosilicate, 
272; nitrate, 167; decomposition by 
heat, 565; oxide, 167; salts, reaction 
with sulfates, 381; sulfate, 167; 
sulfide and chloric acid, 354; and 
hydrogen peroxide, 347 

Lead pencils, 630 

Leather, artificial, 694 

Legumes, assimilation of nitrogen, 515 

Levulose, 682 

Light, ether wave hypothesis, 788 

Lightning, cause of, 773 

Lime, slaking of, 150 

Limelight, 301 

Limestone, 150 

Limewater, 150; test for carbon 
dioxide, 18, 151 

Liquid: equilibrium with vapor, 201; 
state, 198; supercooling of, 206 



Index 



60 1 



[References are to sections, not pages] 



Litharge, 167 

Lithium, 823; flame, 784; spectrum of, 

786 
Litmus: reaction to acids, 89; reaction 

to bases, 88; sensitiveness, 440 
Lubricating oils, 643 
Luminosity, cause of, in flames, 782 
Luncheon, example, 688 
Lye, 162 
Lynde, 776 

Madder, 698 

Magnesia, milk of, 142 

Magnesium: burning of, n, 28, 30, 80; 
in Periodic Table, 832; bromide, 255; 
chloride, 143; fluosilicate, 272; hy- 
droxide, 142, 143, 166; nitrate, 145; 
oxide, 142; sulfate, 144 

Malachite, analysis of, 34 

Maltose, 641, 682 

Manganese and compounds, 342, 837 

Manganese dioxide: catalyst, 306, 320; 
oxidizing agent, 326; preparation of 
chlorine, 234; of bromine, 258 

Maple sugar, 682 

Marble, 150 

Marine acid, 305 

Marsh gas, 643 

Marsh's test, 810 

Matches, 586 

Matter: change of form with tempera- 
ture, 9; conservation, law of, 21; 
electrical nature of, 482, 880; forms 
of, 2 

Meat, 685 

Medicinals, manufacture related to dye 
industry, 699 

Meker burner, 781 

Membrane, semipermeable, 710 

Mendelejeff, 843 ' 

Mercuric compounds, 333; chloride, 
178; nitrate, 178; basic, 564; oxide, 
73, 178, 181; red ash of mercury, 30; 
sulfate, 179 

Mercurous compounds, 2>2>y, chloride, 
180, 182; nitrate, 180; oxide, 181; 
sulfate, 180 

Mercury, 178; ash, decomposition by 
heat, 14; and chlorine, 179; oxides, 
181; and place in Periodic Table, 832; 
properties, 178-82, 832; from red 
oxide, 14; bichloride, 178 

Mesothorium, 870 

Metaboric acid, 802 

Metals, 35; cathions, 390; displace- 
ment by one another, 490; oxidation 
and reduction of, 504; place in 
Periodic Table, 841 



Metaphosphoric acid, 589, 590; and 
albumen, 597 

Metastannic acid, 809 

Metathesis, 383 

Methane: analysis, 55; properties and 
composition, 54 

Methyl compounds: acetate, 657; 
alcohol, 645; ammonium iodide, 658; 
chloride, 644; ether, 646; structural 
formula of, 649; iodide, 660; naph- 
thalene, 669 

Methyl orange, 137, 440 

Methyl violet, adsorption from solution, 
732 

Methylene chloride, 644 

Meyer, Lothar, 844 

Mica, 807 

Micoderma aceti, 653 

Microcosmic salt, 596 

Milk sugar, 682 

Millikan, 467 

Minimum and multiple weights, ex- 
planation of law, 212 

Minimum weights, 63, 64, 212 

Mitscherlich, 810 

Mixed metal, 820 

Moissan, 267, 630 

Moisture, in air, 766 

Molecular heats, 792-93 

Molecular hypothesis, 186, 208 

Molecular weight, 217; and depression 
of the freezing-point, 718; of elec- 
trolytes in solution, 720; and eleva- 
tion of the boiling-point, 718; and 
lowering of vapor pressure of solu- 
tion, 718; from osmotic pressure, 
data, 717; relation to formula weight, 
217 

Molecules, 184, 185; attraction of, 202; 
colors of, in solution, 396; kinetic 
energy of, 198; motion of, 187; per- 
ception of, 703; reality of, 702; 
solubility of, 445; velocity of, 197; 
velocity of, and temperature, 189 

Molybdenite, 813 

Molybdenum: and compounds, 813; 
place in Periodic Table, 836 

Molybdic acid, 813 

Monatomic gases, ratio of molecular 
heats in, 793 • 

Monazite, 820 

Mond process for nickel ores, 816 

Monel metal, 816 

Mordant, 809 

Morse, 712 

Moseley, 848, 87S 

Moth balls, 669 

Mother of vinegar, 653 



602 



Introduction to General Chemistry 



[References are to sections, not pages] 



Multiple weights, law of, 212 
Muriatic acid, 250 
Mustard gas, 695 

Naphthalene, 669 

Natural gas, 643 

Nature of matter, 880 

Neodymium, 820 

Neon, 798; and its family, 825 

Neutrality, of solutions, 433 

Neutralization: acid and base, 89, 379; 
change of conductivity during, 423; 
heat of, 362, 439; ionic theory 
of, 421, 423; of nitric acid by am- 
monium hydroxide, 105; simplified 
equation of, 422 

Newlands, 844 

Newton, first law of motion, 187 

Nickel: and compounds, .816, 817; 
place in Periodic Table, 838 

Nickel ammonium ion, 538 

Nickel coin, 816 

Nickel steel, 816 

Nilson, 845 

Niobium, place in Periodic Table, 835 

Niton, 825 

Nitrates: properties of, 564, 565; 

source of, 540; test for, 549 
.Nitric acid: from air, 567; and 
aluminum hydroxide, 175; from 
ammonia, 570; and ammonium 
hydroxide, 105; anhydride, 555; 
chemical solvent for salts, 560; and 
copper oxide, 165; and magnesium 
hydroxide, 145; and magnesium 
oxide, 145; and metals, 558; from 
nitric oxide, 569; and non-metals, 
559; oxidation by, 557; oxidizing 
agent, 542; properties of, 104, 541; 
from sulfuric acid and niter, 104; 
test for, 549; uses of, 571 

Nitric oxide: combustion in, 546; 
conversion into nitric acid, 569; 
critical temperature, 775; equilibrium 
with nitrogen and oxygen, 566; and 
ferrous sulfate, 548; from ferrous 
sulfate and nitric acid, 547; prepara- 
tion of, 543; properties of, 545 

Nitrides, 514 

Nitrites, 553 

Nitrobenzene, 673 

Nitrocellulose, 693 

Nitro compounds, aromatic, 673 

Nitrogen: assimilation by plants, 515; 
critical temperature, 775; cycle in 
nature, 574; discovery of, 512; 
fixation of atmospheric, <?7«;: and 



hydrogen, 



librium, 520-25; inert part of air, 
15; liquid, from liquid air, 777; 
minimum weight, 63; occurrence of, 
511; and oxygen, 544; and oxygen, 
equilibrium, 566; place in Periodic 
Table, 835; preparation of, 513; 
properties of, 514; pentoxide, 555; 
tetroxide, 545; physical properties 
of 3 55 1 , 55 2 ; preparation of, 550; 
two forms of, 552; trioxide, 554 

Nitroglycerine, 692; adsorption in 
infusorial earth, 726 

Nitrolime, 514 

Nitrophenol, 673 

Nitrosyl chloride, 562, 563 

Nitrosyl sulfuric acid, 616 

Nitro toluene, ortho, meta, and para, 

673 

Nitrous acid, 553; anhydride, 554 

Nitrous oxide, 556; critical tempera- 
ture, 775 

Nonconductor of electricity, 473 

Non-metals, 35; displacement by one 
another, 488; oxidation and reduc- 
tion of, 505; place in Periodic Table, 
841 

Normal solutions, 136-38; 595 

Noyes, A. A., 722 

Nutrition, 688 

Oil: animal and vegetable, 657; of 
bitter almonds, 670, 697; essential, 
697; hardening of, 681; hydrogena- 
tion of, 681; mineral, 643; of vitriol, 

93 
Oildag, 750 
Oleic acid, 680 
Olein, 680 
Oleum, 613 
Onyx, 806 
Opal, 804, 806 

Organic materials, sources of, 701 
Orthoclase, 807 
Orthophosphoric acid. See Phosphoric 

acid, ortho 
Osmic acid, 838 
Osmium, 818, 838; oxide, 838 
Osmosis, 711; in nature, 724 
Osmotic pressure: measurement, 711; 

of sugar solution, 712; theory of, 713, 

7i4 
Ostwald, Wolfgang, 746 
Oxalic acid, 665 
Oxidation, 325; change of valence, 331, 

332, 334, 336; intensity of, 338 
Oxidation and reduction, 501-5; cells, 

502; electronic explanation, 501; by 



and hydrogen, equi- means of electric current. 



Index 



603 



[References are to sections, not pages] 



method of balancing equations, 561; 
potentials, 506 

Oxides: graphic formula of, 324; of 
metallic elements, 313; of non- 
metallic elements, 313 

Oxyacetylene torch, 315 

Oxygen: constancy of concentration in 
air, 766; critical temperature, 775; 
discovery, 305; liquid, 309, 312, 777; 
minimum weight, 63; in nature, 305; 
and nitrogen, 544; and nitrogen, 
equilibrium, 566; in Periodic Table, 
836; from plants, 311; preparation 
of, 306-10; properties of, 16, 312, 313; 
uses of, 315 

Oxone, 307 

Ozone, 316; as a germicide, 317 

Paint: drying of, 364; luminous, 859 

Paintings, restoration of, 347 

Palladium, 838 

Palmitic acid, 678 

Palmitin, 679, 680 

Paraffine, 643; series, 643 

Pentathionic acid, 628 

Perchlorates, 355 

Perchloric acid, 355 

Perfume, 697 

Period of radioactive substance, 864 

Periodic law, 843, 844 

Periodic system, 822-50 

Periodic Table, 829; anomalies of, 848; 
arrangement of Harkins', 849; history 
of, 844 

Permanganates, 343; color of, in solu- 
tion, 39s 

Permonosulfuric acid, 628 

Peroxides, 322; graphic formula of, 324 

Perrin, 478, 703, 707 

Persulfuric acid, 628 

Petroleum, 643 

Phenol, 672 

Phenolphthalein, 137, 440 

Phosgene, 695 

Phosphate rock, 160, 580 

Phosphates: precipitation of, 452; 
qualitative test for, 597; and silver 
nitrate, 171, 597; use and production, 
598 

Phosphine, 588 

Phosphonium chloride, 588 

Phosphorescence, 584; caused by 
X-rays, 852 

Phosphoric acid, ortho: and albumin, 
597; failure to precipitate phos- 
phates, 452; from hydrolysis of 
phosphorus pen tachloride, 247; ioni- 
zation of, 592; normal solution of, 



595; preparation, 158, 589, 590; 
properties, 159, 590-92, 597; and 
silver nitrate, 597; titration of, 595. 
See also Metaphosphoric acid; Pyro- 
phosphoric acid 

Phosphorous acid, 588; from hydrolysis 
of phosphorus trichloride, 247 

Phosphorus: amorphous, 581; and bro- 
mine, 256; and chlorine, 247; dis- 
covery of, 577; and iodine, 264; 
manufacture of, 580; place in Peri- 
odic Table, 835; red, 579, 581; slow 
oxidation of, 583; white, danger in 
use of, 585; white, physical proper- 
ties of, 578; yellow, 578; com- 
pounds, 588; pentabromide, 256; 
pentachloride, 247; pentoxide, 247, 
588; sulfide, 586; tribromide, 256; 
trichloride, 247; tri-iodide, 264; tri- 
oxide, 588 

Photosynthesis, 691 

Phthallic acid, 698 

Picric acid, 673, 693 

Piperonal, 697 

Pitchblende, 874 

Plaster of Paris, 155 

Platinum: adsorption of hydrogen and 
oxygen, 731; and aqua regia, 562; 
and compounds, 818; properties and 
place in Periodic Table, 838 

Polymer, 613 

Potash, 162; caustic, 106; monopoly of 
Germany, 353. See also Potassium 
hydroxide 

Potassium, 106; action on water, 106; 
aluminum sulfate, 175; cupric 
chloride, 175; flame, 784; place in 
Periodic Table, 831, 848; acid tel- 
lurate, 815; antimonyl tartrate, 811; 
bicarbonate, 162; bichromate {see 
Dichromate); bromide, 255; use 
of, 260; carbonate, 162; chlorate: 
oxidizing agent, 326; preparation of 
oxygen from, 306; preparation and 
use, 352; precipitation of, 447; 
chromate, 346; dichromate, 346; and 
hydrogen peroxide, 321; and hydro- 
gen sulfide, 346, 609; and sulfurous 
acid, 346; ferricyanide, 665; ferro- 
cyanide, 665; hydrogen tartrate, 665; 
hydroxide, 106; iodide, and ferric 
chloride, electronic explanation of, 
503; nitrite, from potassium nitrate, 
305; perchlorate, 355; preparation 
of oxygen, 306; permanganate and 
ferrous sulfate, 343; and hydrogen 
chloride, 235, 343; and hydrogen 
sulfide, 609; oxidizing agent, 326, 343; 



604 



Introduction to General Chemistry 



[References are to sections, not pages] 



preparation from potassium man- 
ganate, 837; and sulfurous acid, 343; 
platinic chloride, 818; tellurite, 815; 
uranyl sulfate, rays from, 852 

Potato starch, preparation of, 636 

Potential energy, 368 

Potentials, oxidation-reduction, 506 

Powder: black, 571, 572; smokeless, 
692 

Precipitates, relative stability of differ- 
ent forms, 762 

Precipitation, 154; cause of, 446; classi- 
fication of, 452; conditions favoring, 
454; of electrolytes, 442; by sub- 
stance having a common ion, 453 

Pressure: critical, 775; effect of, on 
gases, 4; effect on system in equi- 
librium, 287; partial, 113 

Priestley, 305 

Proof spirit, 641 

Propane, 58, 643 

Proportions, law of definite, 46 

Protecting agents for colloids, 745 

Protein, 685 

Prout, hypothesis, 851 

Prussic acid, 665 

Ptyalin, 682 

Pure substance, 23 

Pyrite, 610 

Pyrophoric alloy, 820 

Pyrophosphoric acid, 590 

Pyrosulfates, 624 

Pyroxaline, 694 

Quartz, sand, 804 
Quartzite, 806 
Quicklime, 150 

Radical, 147; acid, 377; acid, anions, 
391; basic, 378; phenyl, 670 

Radioactive change, theory of, 865 

Radioactive substances: atomic weights 
of, 869; arrangement of, in Periodic 
Table, 873; valence of, 872 

Radiothorium, 870 

Radium: discovery of, 855-56; emana- 
tion, 863, 864; heat from, 862; how 
measured and sold, 875; origin of, 
867; place in Periodic Table, 845, 
857; rays of, 480; series, 866; 
technical production, 874; total 
world-production, 874; use in thera- 
peutics, 876 

Radium clock, 860 

Rain, formation of, 772 

Ramsay, 791, 845 

Raoult, 718, 719 



Rare earth elements, 820; place in 

Periodic Table, 846 
Rayleigh, 791 
Rays of radium, 480, 858-61. See also 

X-rays 
Reaction velocity, kinetic hypothesis 

applied to, 275-78 
Reducing agents, intensity of, 338 
Reduction, 327; and change of valence, 

335> 336; electronic explanation of, 

501; of metals, 504; of non-metals, 

505 
Refractory substance, 568 
Respiration, 314 
Rhodium, 838 
Rock crystal, 806 
Rock salt, 24 
Roentgen rays, 476 
Rose's metal, 812 
Rowland, 468 
Rubber: a colloid, 760; synthetic, 700; 

vulcanization of, 604 
Rubidium, 823 
Ruby, 174 

Ruthenium, place in Periodic Table, 838 
Rutherford, D., 512 
Rutherford, E., 481, 865 

Sal soda, 26 

Salicylic acid, 697 

Safrole, 697 

Salt, common: composition, 45; elec- 
trolysis of, 238; preparation, 41; 
purification from rock salt, 24 

Salt, parts of, 376 

Saltpeter, Chile, 104; occurrence, 540 

Salts: acid, 102; active parts of, 376; 
a class of chemical substances, 92; 
hydrolysis of, 436; ionization of, 
410; little soluble of strong acids, 
not dissolved by acids, 458; neutral, 
102; primary, secondary, and terti- 
ary, 159 

Sandpaper, 806 

Sandstone, 806 , 

Saponification, 679 

Sapphire, 174 

Sassafras, oil of, 697 

Scandium, discovery of, 845 

Scheele, 305 

Scurvy, 689 

Seaweeds, iodine from, 261 

Selenium, and its compounds, 815, 836 

Series: displacement of metals, 492; 
of non-metals, 488 

Siedentopf, 706 

Silica, 806 



Index 



605 



[References are to sections, not pages] 



Silicates, 807 

Silicic acid, 807; a gel, 761 

Silico-chloroform, 805 

Silico-ethane, 805 

Silico-mesoxalic acid, 805 

Silico-methane, 805 

Silico-oxalic acid, 805 

Silicon: . compared to carbon com- 
pounds, 805; and compounds, 804-8, 
827; place in Periodic Table, 834; 
dioxide, 806; fluoride, 270; hexa- 
chloride, 805; tetrachloride, 805 

Silk, artificial, 694 

Silver, 168; colloidal, 744; place in 
Periodic Table, 831; test for, 169; 
acetate: precipitation, 452; solution 
by nitric acid, 456; ammonium ion, 
533S4; effect of acids on, 536; 
ammonium salts, 535, 537; azid, 
531; bromide, 257; use, 260; chloride, 
169; solubility in ammonia solution, 
532-35; hydroxide, 172; nitrate, 168; 
electrolysis of solution, 387; and 
hydrochloric acid in the presence of 
gelatine, 745; oxide, 172; phosphate, 
171, 597; salts, reaction with chloride, 
382; sulfate, 170; sulfide, precipita- 
tion of, 452; thiosulfate, complex 
ion, 626 

Smoke screens, 587 

Soap: cleansing action, 751-54; com- 
position, 162, 678 
'Soda, 26; caustic, and hydrochloric 
acid, 89; caustic, from sodium and 
water, 40; washing, 26 

Sodamide, 527 

Soda water, 633 

Soddy, 481, 865 

Sodium: and chlorine, theory of union 
of, 485; flame, 784; place in Periodic 
Table, 831; spectrum of, 786; and 
water, 40; acetate, 157; action of 
water on, 435; and hydrochloric 
acid (ionic theory), 424; aluminate, 
177; aluminum fluoride, 267; alu- 
minum sulfate, 175; antimony flu- 
oride, 811; benzoate, 671; bicar- 
bonate, 161; bromide, 255; use of, 
260; carbonate, 161; chlorate, 353; 
chloride, action on sulfuric acid, 460; 
electrolysis of solution, 385; pre- 
cipitated by hydrochloric acid, 453; 
fluoride, 269, 270; fluosilicate, 272; 
hydrogen sulfate, 98, 100; action of 
hydrochloric acid on, 289; hydro- 
sulfite, 628; hydroxide, 88; and 
acetic acid (ionic theory), 434; and 



ammonium chloride, 426, 462; and 
hydrochloric acid, 421; hypochlorite, 
350; metaphosphate, 594; nitrate, 
104; decomposition by heat, 565; 
nitrite, 513; oleate, work of Harkins 
on, 754; perchlorate, 355; peroxide, 
307; phosphates, 159, 594; pyrophos- 
phate, 594; silicate, 270, 807; silver 
thiosulfate, 626; stannite, 809; sul- 
fate, 94; hydrate of, 94, 96, 97; 
sulfides, 607; tetraborate, general 
properties, 801, 803; tetrathionate, 
625; thiosulfate, 625; tungstate, 814; 
zincate, 177 

Solder, 809 

Solid state, theory of, 203 

Solubility: curves, 132; effect of crys- 
talline form on, 133; effect of tem- 
perature on, 132-34; gases in liquids, 
125; and heat of solution, 134; 
liquids in liquids, 124; of molecules, 
445; of solids, 122; two kinds, 135 

Solution, 120; compared to gas, 708; 
comparison with suspension, 708; 
concentration of, 121; electrolysis 
of, 385; heat of, 127; kinetic theory 
of, 443; of little soluble electrolytes, 
455; normal, 136-38, 595; saturated, 
122; supersaturated, 123; Van't 
Hoff's theory, 715 

Soot, formation in flame, 782 

Spark spectra, 786 

Specific gravity, no 

Specific heat, in; and symbol weight, 
229 

Spectra: bright-line, 786; continuous, 
793; dark-line, 789; of gases and 
volatile substances, 787 

Spectroscope, 785 

Spectroscopic analysis, 786 

Spelter, 148 

Spinthariscope, 859 

Spirit, proof, 641 

Stannic compounds, 809 

Stannous compounds, 809 

Stannum, 809 

Starch: glucose from, 636; occurrence 
of, 636; properties of, 637, 638; test 
for, 637 

Steam, 8; action on iron (chemical 
equilibrium), 290; burns by, 117; 
production, fuel for, 360; use for 
heating, 116 

Stearic acid, 678 

Stearin, 680 

Stibine, 811 

Stoney, 466 



6o6 



Introduction to General Chemistry 



[References are to sections, not pages] 



Strontium, 832; flame, 784 

Strutt, 860 

Sublimation, 179 

Subscripts, chemical, 61 

Substance, 22, 23; properties of a, 108 

Sugars, structure of, 683 

Sulfates, 624; test for, 164, 380 

Sulfides: precipitation of, 608; pre- 
cipitation by ammonium sulfide, 452 

Sulfites, 612 

Sulfocyanates, 665 

Sulfur: commercial importance of, 604; 
consumption of, 604; flowers of, 
603; heat of combustion, 357; 
monoclinic, 600; place in Periodic 
Table, 836; plastic, 600; properties 
of, 600, 601; rhombic, 600; rock, 
603; roll, 603; sources of, 602; 
dioxide: critical temperature, 775; 
preparation and properties, 610; 
and water, equilibrium, 286; mono- 
chloride, 600; trioxide, 613 

Sulfuric acid: and aluminum hydroxide, 
175; anhydride of, 613; and copper 
oxide, 165; dilute, action on iron, 
173; dilute, action on magnesium, 
149; dilute, action on zinc, 149; fum- 
ing, 613; importance of, 614; and 
magnesium hydroxide, 144; and 
magnesium oxide, 145; manufacture 
of, 615; neutralization of, 94; prop- 
erties, 92, 620-23; and sodium 
chloride (ionic theory), 460 

Sulfurous acid, 340, 611; equilibrium, 
sulfur dioxide and water, 286; and 
iodic acid, 277 

Sun: composition of, 790; source of 
energy, 374 

Supercooling, 206 

Superphosphate, 160 

Surface, increase with subdivision, 727 

Suspensions, 120; laws governing, 707 

Suspensoids, 735, 747; of commercial 
importance, 750; influence of the 
charge on stability of, 740 

Symbol weights: accuracy of , 2 2 2 ; how 
found, 224, 228; oxygen basis, 223; 
product of, and specific, heat, 229 

Symbols: chemical, 61; and minimum 
weights, 66 

Synthesis, 31; organic, 696, 698 

Table sugar, 682 
Talc, 807 

Tannin, used in making red gold solu- 
tions, 746 
Tantalum, 835 



Tantiron, 804 

Tartar emetic, 811 

Tartaric acid, 665 

Tellurium and its compounds, 815, 836, 
848 

Temperature: absolute, 6; critical, 
775; effect of change of, on equi- 
librium, 288; effect of change of, on 
gases, 5 ; effect of change of, on solu- 
bility, 134, 288 

Tetramethyl silico-methane, 805 

Tetrathionic acid, 628 

Thermal Unit, British, 358 

Thermite, 330 

Thermometer: centigrade, 6; gas, 6 

Thermos bottles, 777 

Thiocyanate (sulfocyanate), 665; and 
ferric chloride, 280, 415 

Thorium: and its active products, 870; 
and its compounds, 820; in monazite, 
820; place in Periodic Table, 834; 
properties, etc., 834 

Thorium oxide, use in gas mantles, 301 

Thorium X, 870 

Tin: and its compounds, 809, 834; 
amalgam, foil, plate, ware, 809; 
place in Periodic Table, 834 

Titanium, 834 

Titration, 137; indicators for, 440 

T.N.T., 673, 693 

Transition point, 600 

Transmutation of elements, 851 

Trimethyl amine, 59, 658 

Tri nitro phenol, 673, 693 

Tri nitro toluene, 673, 693 

Tripoli, 806 

Trithionic acid, 628 

Tungsten: and its compounds, 814, 
836; use in incandescent filaments, 
814 

Turkey red, 698 

Turpentine, 248 

Tw. See Twadell 

Twadell, scale of specific gravity, 618 

Tyndall, J., 748; effect, 748 

Type metal, 811 

Ultramicroscope, 703, 706 

Unit volume, the chemical, 65 

Univalent, 146 

Unsaturated compounds, 527, 663 

Uranium: parent of radium, 867; 

properties and compounds, 836; 

series, 868 
Uranyl nitrate, 836 
Urea, 659; synthesis of, 696 



Index 



607 



[References are to sections, not pages] 



Valence, 146; atoms and radicals, 183; 
change of, in oxidation and reduction, 
331-36; electrical nature of, 484, 
880; and Periodic Table, 850; posi- 
tive and negative, 840; of radicals, 
147; and the structure of inorganic 
molecules, 839; table of, 183 

Valeric acid, 680 

Vanadium, 835 

Vanilla, 697 

Vanillin, 697 

Van't Hoflf, 715, 719 

Vapor, equilibrium with liquid, 201 

Vapor pressure: correction of volume 
of gas for, 114; of hydrates, 131; 
liquids and solids, 114; lowering of, 
by dissolved substance, 129; theory 
of, 200; of water, 112 

Varnish, drying of, 364 

Vaseline, 643 

Velocity of reactions, 275 

Vinegar, 157; cider, 653; plant, 653 

Vitamines, 689 

Vitriol: blue, 497; oil of, 619 

Volt, 499 

Volume, atomic, 823 

Warfare, materials for, chemical, 695 

Water: adsorption by glass, 730; com- 
position of, 36; distilled, 23; elec- 
trolysis of, 27, 295; formation of, 
139; formula of, 70; hard, 156, 803; 
ionization of, 420; a product of 
combustion, 20; properties, 108-12; 
reactions of, 140; and sodium, 40; 
synthesis of, 37; vapor, in air, 766; 
vapor pressure, 112 

Water glass, 808 

Water soluble B, 689 



Weight, equivalent, 403 

Weights: atomic, 800 (see also inside 
of back cover) ; law of minimum and 
multiple, 64, 212; minimum of ele- 
ments, 63; minimum and symbols, 
66; symbol and formula weights 74 

Whetstone, 806 

Whiskey, 640 

Willemite, 807; phosphorescent, 861 

Wine, 640 

Winkler, 845 

Wintergreen, oil of, 697 

Wohler, 696 

Wollaston, 790 

Wollastonite, 807 

Wood alcohol, 645 

Wood's metal, 812 

Work, 368 

Wulframite, 814 

Xenon: discovery and properties, 798; 

and family, 825 
X-rays, 476; and crystal structure, 204, 

879; spectra, 877 
Xylene, 668 

Yeast, 640 

Zinc: and concentrated sulfuric acid, 
622; and dilute sulfuric acid, 149; 
and hydrochloric acid, 149; and nitric 
acid, 561; place in Periodic Table, 
832; properties and use, 148; chlo- 
ride, 148; nitrate, 148; oxide, 148; 
sulfate, 148; sulfide: and alpha rays, 
859; precipitation from solution, 452 

Zirconium, and compounds, 834 

Zsigmondy, 706 



PERIODIC TABLE ACCORDING TO ATOMIC NUMBERS, H:i 



Pe- 
riod 





A B 


II 
A B 


III 
A B 


IV 

A B 


V 

A B 


VI 
A B 


VII 
A B 


VIII 


I 


2 

He 


3 
Li 


4 
Be 


5 
B 


6 
C 


7 

N 


8 



9 
F 




2 


IO 

Ne 


11 
Na 


12 
Mg 


13 

Al 


14 
Si 


15 
P 


16 

s 


17 
CI 






18 
A 


19 20 
K Ca 


21 
Sc 


22 
Ti 


23 
V 


24 

Cr 


25 
Mn 


26 27 28 
Fe Co Ni 


3 




29 30 

Cu Zn 


31 
Ga 


32 
Ge 


33 

As 


34 
Se 


35 
Br 






36 
Kr 


37 1 38 
Rb Sr 


39 
Y 


40 
Zr 


41 
Cb 


42 
Mo 


43 
? 


44 45 46 
Ru Rh Pd 


4 




47 48 
Ag Cd 


49 
In 


50 
Sn 


5i 
Sb 


52 
Te 


53 
I 






54 
Xe 


55 56 
Cs iBa 


57 
La 


58 
Ce 


59 
Pr 


60 
Nd 


61 

? 




5 




62 
Sa 


63 
Eu 


64 
Gd 


65 
Tb 


66 
Dy 


67 
Ho 


68 
Er 




6 




69 
Tm 


70 

? 


7i 
Yb 


72 
Lu 


73 
Ta 


74 
W 


75 

? 


76 77 78 
Os Ir Pt 




79 I 80 
Au Hg 


81 
Tl 


82 
Pb 


&3 
Bi 


84 

? 


85 

? 




7 


86 

Nt 


87 

? 


88 
Ra 


89 
? 


90 
Th 


91 

? 


92 

u 







H 4 84 Iil| 








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